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Thermodynamic properties of water + normal alcohols and vapor-liquid equilibria for binary systems of methanol or 2-propanol with water
HglBPgISi
EQUIUBIilA ELSEVIER
FluidPhaseEquilibria127 (1997) 181-190
Thermodynamic properties of water + normal alcohols and vapor-liquid equilibria for binary systems of methanol or 2-propanol with water Brahim Khalfaoui a, *, Abdessalem H. Meniai b, Rafael B o r j a c a Chemical Engineering Department, UMIST, Sackville Street, Manchester, M60 1QD, UK b Institut de Chimie Industrielle, Universite de Constantine, 25000 Constantine, Algeria c Instuto de la Grasa y sus Derivados (CSIC), Avda. Padre Garcia Tejero 4, 41012 Seville, Spain
Received 15 January 1996; accepted 17 May 1996
Abstract
Experiments have been carried out to determine densities and refractive indices of binary and ternary mixtures of water with normal alcohols (methanol and n-butanol) at 293.15 K. Also, vapor-liquid equilibrium data have been measured for binary systems of methanol-water and 2-propanol-water at atmospheric pressure, from which the excess Gibbs free energy was computed for both systems. The experimental vapor-liquid equilibrium (VLE) data have been compared with those predicted by the Wilson, NRTL and UNIQUAC equations. Agreement has been found to be very satisfactory. Keywords: Methanol-H20 systems; 2-Propanol-H20 systems; Density; Refractive index; Gibbs free energy; Vapor-liquid equilibria; Water-n-alcohol systems
I. Introduction
Vapor-liquid equilibrium data are always required for engineering use such as in design and, especially, in distillation, which is the most common operation performed in the chemical industry for the separation of liquid mixtures. In previous work [1,2] we presented experimental results of vapor-liquid equilibrium (VLE) measurements for binary mixtures of water with chlorinated hydrocarbons and water with brominated hydrocarbons. The present work reports the VLE results of measurements on binary mixtures of methanol-water and 2-propanol-water at atmospheric pressure. The experimental results have been compared with those predicted by the Wilson [3], NRTL
* Corresponding author (on leave from Universite de Constantine, 25000 Constantine, Algeria). 0378-3812/97/$17.00 Copyright© 1997ElsevierScienceB.V. All rightsreserved. PII S0378-3812(96)031 29-9
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190
182
(non-random, two-liquid) [4] and UNIQUAC [5] equations. The methanol-water system has been investigated previously [6-9] and this system is, therefore, suitable for testing the performance of the VLE apparatus used in this work. However, new reliable data were obtained for this system. Additionally, new accurate and reliable VLE data for the 2-propanol-water system are also obtained in the present study and are compared with those of Wilson and Simon [10], Choffe et al. [11], Verhoeye [12], and Verhoeye and De Schepper [9]. Prior to the VLE measurements, the densities and refractive indices of methanol-water, n-butanol-water, methanol-n-butanol and of the ternary system methanol-n-butanol-water were measured at 293.15 K. This was done as a continuation of a long-term study on the thermodynamic properties of liquid mixtures containing water and normal alcohols in general. Those measurements of particular interest have been made in order to expand the range of mixtures for which the physical properties can be calculated. Furthermore, VLE measurements on binary and ternary mixtures composed of methanol, n-butanol and water were previously carried out [ 13]. The VLE results of this work were used to compute the excess Gibbs free energy for the two studied systems.
2. Experimental section 2.1. Materials Methanol and n-butanol were supplied by Fisons Scientific Equipment Ltd., and 2-propanol was supplied by BDH Chemicals Ltd.. The purity of these alcohols was checked by gas chromatography (GC) and was found to be in excess of 99%. The alcohols were used without further purification. The water used was doubly distilled tap water. Table 1 compares some of the measured properties with literature values. As can be seen, the agreement is good.
2.2. Apparatus and procedure The densities of the pure liquids and the mixtures were measured at 293.15 K using an Anton Paar DMA 46 densimeter with a precision of _0.0001 g cm -3. The refractive index measurements at 293.15 K were carried out using a dipping refractometer which was equipped with thermoprisms (Carl Zeiss, Jena) that had a measurement precision of + 2 × 10 -5. Illumination was provided by a sodium lamp. Both instruments were statically controlled at 293.15 + 0.05 K by circulating water from a
Table 1 Densities p, refractive indices r/ and normal boiling points Tb of the pure materials used in this study at 293.15 K Component p (g cm -3) 7/(D) Tb (K) Methanol 2-Propanol n-Butanol Water
Exptl. 0.7915 0.7853 0.8100 0.9980
Lit. 0.7915 [18] 0.78534 [ 16] 0.8098 [ 17] 0.9981 [20]
References are given in square brackets.
Exptl. 1.32900 1.37780 1.39913 1.33301
Lit. 1.32885 [19] 1.37745 [ 16] 1.39920 [ 17] 1.33320 [20]
Expti. 337.80 355.50 390.80 373.13
Lit. 337.68 [ 18] 355.41 [18] 390.15 [15] 373.20 [21]
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
183
thermostat. The VLE measurements were carried out using a vapor-liquid equilibrium apparatus. This consisted of an all-glass dynamic recirculating still described by Walas [ 14], equipped with a Cotterell pump. The apparatus used (Labodest model, manufactured by Fischer Labor und Verfahrenstechnic, Germany) is designed for a charge of approximately 100 cm 3 of mixture and can be operated at low, moderate or high pressures, ranging from 0.25 kPa through atmospheric pressure up to 400 kPa, and at temperatures up to 523.15 K. The vapor and liquid phases reach equilibrium once the vapor temperature has remained constant for a period of 15-30 min. At this moment, the first sample take-off can begin. The equilibrium temperature was measured with a digital Fischer thermometer with a precision of 0.01 K, and the pressure with a digital manometer with a precision of 0.01 kPa. For further details, the apparatus used in this work was described elsewhere by Aucejo et al. [15]. 2.3. Analysis The liquid and the condensed vapor compositions were determined using a Hewlett Packard 5890 S II gas chromatograph equipped with a thermal conductivity detector (TCD). The GC response peaks were integrated by using an electronic integrator. The packed chromatographic column used was a fused silica capillary column 10 m in length and 0.35 mm in internal diameter, coated with Poraplot Q. Previous calibration analysis was carried out with mixtures of known composition. Usually, a single analysis of the vapor or liquid composition by GC is imprecise. However, with repeated measurements, the standard deviation of a composition analysis was generally less than 0.001 mole fraction. At least two analyses were performed for each liquid and each vapor composition except for the methanol-water system, whose compositions were determined from both GC and density measurements in order to check the reproducibility of the analytical results.
3. Results and discussion
The physical properties of the pure components used in this work are given in Table 1 along with the literature values taken from Refs. [15-21]. The experimental densities and refractive indices of the three binary systems and the ternary system at 293.15 K are given in Table 2. The densities were measured with an uncertainty of _ 0.0002 g cm 3. The experimental VLE results for the two binaries, methanol-water and 2-propanol-water at atmospheric pressure together with the calculated values for the liquid phase activity coefficients (3,1 and 3'2) and the excess Gibbs free energy are given in Tables 3 and 4, respectively. The activity coefficients, 3'i, in the liquid phase were calculated using the following equation 3"i = Y i P / x i P°
(1)
where the vapor pressures pO were calculated with the Antoine equation given in Table 5 along with the Antoine constants taken from Ref. [22] by assuming ideal behavior of the mixture in the vapor phase. The Yi values computed using Eq. (1) are listed in Tables 3 and 4. It can be seen that both systems present a positive deviation from ideality. A comparison of the VLE results for the methanol-water system with the data cited earlier in this paper is made in Fig. 1, where the vapor phase composition is plotted against the liquid phase composition. As can be seen, the agreement is generally good considering the difficulties inherent in measuring vapor-liquid equi!ibria in the dilute
B. Khalfaoui et al./Fluid Phase Equilibria 127 (1997) 181-190
184
Table 2 Densities and refractive indices of solutions of water (1), methanol (2) and n-butanol (3) at 293.15 K W1
100 0.0 0.0 6.37 12.43 18.30 29.91 64.65 7.15 12.62 32.69 46.08 79.59 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 6.50 8.30 8.60 8.90 9.00 11.20 13.03 19.23 20.64 25.60 26.73 29.44 36.53 36.71 52.45 60.66 71.41
W2
0.0 100 0.0 93.63 87.57 81.70 70.09 35.35 0.0 0.0 0.0 0.0 0.0 92.15 74.50 60.19 14.35 12.75 7.45 2.75 1.43 81.20 17.10 35.20 52.30 84.90 67.60 32.57 49.52 38.83 33.22 53.07 22.65 52.44 41.07 34.97 21.49 6.56
W3
0.0 0.0 100 0.0 0.0 0.0 0.0 0.0 92.85 87.38 67.31 53.92 20.41 7.85 25.50 39.81 85.65 87.25 92.55 97.25 98.57 12.30 74.60 56.20 38.80 6.10 21.20 54.40 31.25 40.53 41.18 20.20 47.91 11.03 22.22 12.58 17.85 22.03
P
(g cm- 3)
(D)
0.9980 0.7915 0.8100 0.8201 0.8328 0.8430 0.86326 0.9350 0.8307 0.8417 0.8787 0.9066 0.9848 0.7971 0.7981 0.7990 0.8074 0.8076 0.8086 0.8096 0.8099 0.8111 0.8224 0.8197 0.8170 0.8112 0.8213 0.8347 0.8404 0.8492 0.8610 0.8659 0.8679 0.8876 0.8850 0.9181 0.9319 0.9665
1.3330 1.3290 1.3991 1.3298 1.3300 1.33 I0 1.3314 1.3320 1.3941 1.3911 1.3606 1.3501 1.3416 1.3405 1.3486 1.3568 1.3890 1.3902 1.3942 1.3982 1.3986 1.3401 1.3822 1.3710 1.3578 1.3336 1.3476 1.3725 1.3538 1.3659 1.3681 1.3490 1.3720 1.3490 1.3568 1.3563 1.3540 1.3438
r e g i o n s o f m i x t u r e s . T h i s w a s d o n e to e v a l u a t e the p e r f o r m a n c e o f t h e V L E a p p a r a t u s u s e d in this w o r k . T h e V L E r e s u l t s o f t h e 2 - p r o p a n o l - w a t e r s y s t e m , set o u t in T a b l e 4, a r e a l s o p r e s e n t e d in g r a p h i c a l f o r m as a p l o t o f v a p o r p h a s e c o m p o s i t i o n v e r s u s l i q u i d p h a s e c o m p o s i t i o n in F i g . 2, w h e r e the a z e o t r o p i c p o i n t is s i t u a t e d at x = y = 0.689. A s c a n a l s o b e o b s e r v e d , the V L E r e s u l t s o b t a i n e d in this s t u d y a n d s h o w n o n the p l o t a r e in g o o d a g r e e m e n t w i t h t h o s e f o u n d in the l i t e r a t u r e c i t e d in the
185
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
Table 3 Vapor-liquid equilibrium data, pressure P, temperature T, liquid phase mole fraction x~, vapor phase mole fraction Yl, activity coefficients Yi and excess Gibbs free energy G E / R T , for the methanol (1) + water (2) system P (kPa)
T (K)
x1
Yl
3Jl
T2
GE/RT
99.988 99.988 100.042 99.782 100.135 100.155 100.135 100.228 100.682 100.727 100.737 100.737
337.80 337.85 337.90 337.95 340.85 344.65 347.55 350.65 354.55 357.75 362.65 373.15
1 0.8948 0.8150 0.7649 0.6060 0.4709 0.3477 0.2561 0.1757 0.1174 0.0574 0.0
1 0.9498 0.9115 0.9005 0.8309 0.7766 0.7078 0.6472 0.5852 0.4928 0.3742 0.0
1.0428 1.0972 1.1496 1.2002 1.2491 1.3835 1.5348 1.7672 1.9920 2.6174 -
1.9386 1.9401 1.7081 1.5286 1.2756 1.1964 1.11 44 1.0136 1.0196 0.9742 -
0.0 0.1071 0.1982 0.2325 0.2777 0.2335 0.2300 0.1903 0.1112 0.098 0.0306 0.0
Table 4 Vapor-liquid equilibrium data, pressure P, temperature T, liquid phase mole fraction Xl, vapor mole fraction y~, activity coefficients ~/~ and excess Gibbs free energy G E / R T for the 2-propanol (1) + water (2) system P (kPa)
T (K)
xI
Yl
Yl
72
GE/RT
102.27 102.27 102.27 102.12 102.27 102.25 102.13 102.40 102.03 101.95 101.98 101.98
355.5 352.35 352.85 353.15 353.25 353.55 354.05 354.45 356.35 356.95 357.75 373.13
1 0.8495 0.7975 0.7541 0.6628 0.4887 0.3241 0.2407 0.1350 0.0840 0.0456 0.0
1 0.8015 0.7512 0.7201 0.6854 0.6150 0.5641 0.5517 0.4764 0.4582 0.4380 0.0
1.0799 1.0565 1.0565 1.1412 1.3717 1.8575 2.4128 3.4309 5.1743 8.8313 -
2.945 2.6883 2.4568 2.0086 1.6011 1.3421 1.2123 1.1 48 1 1.0950 1.0560 -
0.0 0.2278 0.2440 0.2624 0.3227 0.3951 0.3995 0.3582 0.2859 0.2211 0.1513 0.0
Table 5 Antoine constants of materials used in VLE
a
Component
A
B
C
Methanol 2-Propanol Water
8.07246 7.56634 7.96681
1574.990 1366.142 1668.210
238.860 197.970 228.00
a logloPiO= A - B / ( t + C ) , Centigrade.
where the units of Pi0 are millimeters of mercury and the temperature is in degrees
186
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190
e II:F 06
'~
y
05
,
/
!
i
C4
p
L
//
02
/
/
/
)/
C.O 0.0
]
i 02
0 4
08
06
10
X 1
Fig. 1. Vapor-liquid equilibrium compositions for methanol-water at atmospheric pressure. (O) This work; ([]) Hala et al. [6]; (zx) Dalager [7]; ( 0 ) Maripuri and Ratcliff [8]; ( v ) Verhoeye and Schepper [9].
introduction. Because of this good agreement, it can be concluded that the experimental errors are trivial for both systems. The excess Gibbs free energy was obtained using the following thermodynamic relationship (2)
G S = RT~_+xi ln yi i
I.C
7
r
-
-
T
. . . . .
~
-
-
08
oo
+'+ /
//
/ 0.0 0.0
0.2
I
I
[
0,4
O.6
0.8
X 1
Fig. 2. Vapor-liquid equilibrium compositions for 2-propanol-water at atmospheric pressure. (O) This work; ([q) Wilson and Simon [ 10]; (zx) Choffe et al. [11 ]; (O) Verhoeye [12]; ( v ) Verhoeye and Schepper [9].
B. Khalfaoui et al. /Fluid Phase Equilibria 127 (1997) 181-190
187
0.3
0.2 E--
0.1
i
0.0( 0.0
0.2
I 0.4 X1
I
I
0.6
0.8
1.0
Fig. 3. Excess Gibbs free energy as a function of the liquid mole fraction for the methanol (1) + water (2) system. (O) Experimental points; ( - - ) splined curve.
and the results for 2-propanol-water and m e t h a n o l - w a t e r are shown in Figs. 3 and 4, respectively. As can be seen from these figures, the two systems exhibit a behavior that is slightly unsymmetrical. The excess Gibbs free energy expressed as GE/RT shows a maximum almost at x I = x 2 for the two systems. The fitness of the curves is an indication o f the precision of the data obtained; therefore, it can be seen in Figs. 3 and 4 that good results have been achieved for the two systems. The results have been tested for thermodynamic consistency using the point-to-point method o f Van Ness et al. [23] modified by Fredenslund et al. [24]. A four-parameter Legendre polynomial was
,
i
1
i
0.4
c.3
L~
0
0.2
0.1
0.0
0.0
t
i
J
,
0.2
0.4
O.6
0.8
1.0
x 1
Fig. 4. E x c e s s G i b b s f r e e e n e r g y as a f u n c t i o n o f the l i q u i d m o l e f r a c t i o n f o r the 2 - p r o p a n o l (1) + w a t e r ( 2 ) s y s t e m . (o) E x p e r i m e n t a l points; ( - - ) splined curve.
B. Khalfaoui et a l . / Fluid Phase Equilibria 127 (1997) 181-190
188
Table 6 Parameters of the Wilson, NRTL and UNIQUAC models
Model Wilson NRTL UNIQUAC
Parameters
Methanol (1) + water (2)
(cal mo1-1 )
NRTL (al2 = 0.30)
MAD(y)
NRTL (oil2 ~ 0.45)
MAD(y)
3,12- 3,22 3,21- Al i g 12- g22 g21- gJl
- 714.80 1805.11 334.44 326.99 1462.45 - 539.00
0.010
756.23 1292.24 701.00 1726.71 562.51 - 82.86
0.010
ul2-u22 u21- u ll
2-Propanol (1) + water(2)
0.009 0.008
0.009 0.009
used for the excess G i b b s free energy. According to Fredenslund et al., the P - T - x - y data are consistent if the m e a n absolute deviation between calculated and measured m o l e fractions o f c o m p o n e n t 1 in the v a p o r p h a s e ( M A D ( y ) ) is less than or equal to 0.01. The values o f M A D ( y ) , for both systems, were found to lie within the stated tolerance, as can be seen in T a b l e 6. As indicated by this test, it can be concluded that our results are consistent. T h e activity coefficients for the two binaries investigated in this study, given in Tables 3 and 4, are c o m p a r e d with those predicted b y the Wilson, N R T L and U N I Q U A C equations and are shown in Figs. 5 and 6 as plots o f In 'Yi versus the liquid phase composition. As can be seen f r o m both graphs, the a g r e e m e n t is satisfactory for the three models. Moreover, as can be observed, the change in In 'Yi with c o m p o s i t i o n is an adequate demonstration o f the t h e r m o d y n a m i c consistency o f the obtained results. The adjustable binary
1.2
1 3
I
0.8 L-'N~X 2 ~.~. ,n71
InT2
I
3-~
/y
x1
Fig. 5. Activity coefficients as a function of liquid compositions for the methanol (1) + water (2) system. (O, O) This work. Curve 1, Wilson; curve 2, NRTL; curve 3, UNIQUAC.
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
25~_
2~
2.0 I_~o
t
,
~
1 5
00
189
~
02
0 a xI
0.6
0 g
1.0
Fig. 6. Activity coefficients as a function of liquid compositions for the 2-propanol (1) + water (2) system. ( O , 0 ) This work. Curve 1, Wilson; curve 2, NRTL; curve 3, UNIQUAC.
parameters of each model were estimated in the same manner as performed by Benito et al. [25], by using a non-linear regression method that minimizes the following objective function N
E i= l
[() ")/1
"Yl
In - -
"Y2
-In i,calc
-")/2
(3) i,expt
The resulting parameters from the fit to each model are shown in Table 6.
4. List of symbols A,B,C gij
P
P? R t
T uij
W xi Yi
Antoine constants in Table 5 interaction parameter in NRTL equation given in Table 6 pressure pure component vapor pressure gas constant in Eq. (2) temperature (°C) absolute temperature interaction parameter in UNIQUAC equation given in Table 6 weight percentage in Table 2 liquid phase mole fraction vapor phase mole fraction
4.1. Greek letters Ol-ij
")'i
non-randomness parameter in NRTL equation given in Table 6 activity coefficient
19o
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190 refractive index interaction e n e r g y in W i l s o n e q u a t i o n g i v e n in T a b l e 6 density (g crn - 3 )
Aij P
4.2. Subscripts i, 1,2,3
c o m p o n e n t identification
4.3. S u p e r s c r i p t s o
pure c o m p o n e n t
References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25]
M.R. Cooling, B. Khalfaoui and D.M.T. Newsham, Fluid Phase Equilibria, 81 (1992) 217-229. B. Khalfaoui and D.M.T. Newsham, Fluid Phase Equilibria, 98 (1994) 213-223. G.M. Wilson, J. Am. Chem. Sot., 86 (1964) 127-130. H. Renon and J.M. Prausnitz, AIChE J., 14 (1968) 135-144. D.S. Abrams and J.M. Prausnitz, AIChE J., 21(1) (1975) 116-128. E. Hala, J. Wichterle, J. Polak and T. Boublik, Vapour-Liquid Equilibrium Data at Normal Pressures, Pergamon Press, London, 1968. P. Dalager, J. Chem. Eng. Data, 14 (1969) 298-301. V.O. Maripuri and G.A. Ratcliff, J. Chem. Eng. Data, 17 (1972) 366-369. L. Verhoeye and H.De Schepper, J. Appl. Chem. Biotechnol., 23 (1973) 607-619. A. Wilson and E.L. Simon, Ind. Eng. Chem., 44(2) (1952) 2214-2219. B. Choffe, M. Cliquet and S. Meunier, Rev. Inst. Ft. Pet., 15(6) (1960) 1051-1060. L.A.J. Verhoeye, J. Chem. Eng. Data, 13 (1968) 462-467. D.M.T. Newsham and N. Vahdat, Chem. Eng. J., 13 (1977) 27-31. S.M. Walas, Phase Equilibria in Chemical Engineering, Butterworth, London, 1985. A. Aucejo, M.C. Burguet, J.B. Monton, R. Munoz, M. Sanchotello and M.I. Vazquez, J. Chem. Eng. Data, 39 (1994) 271-274. A. Aucejo, V.G. Alfaro, J.B. Monton and M.I. Vazquez, J. Chem. Eng. Data, 40 (1995) 332-335. A. Weissberger, E.S. Proskauer, J.A. Riddick and E.E. Toops, Jr., Organic Solvent, Interscience, London, 1955. J. Timmermans, Physico-Chemical Constants of Pure Organic Compounds, Vol. 1. Elsevier, New York, 1950. R.C. Weast (Ed)., Handbook of Chemistry and Physics, The Chemical Rubber Co., Cleveland, OH, 1968. E.G. Mahers and R.A. Dawe, J. Chem. Eng. Data, 31 (1986) 28-31. R.C. Reid, J.M. Prausnitz and T.K. Sherwood, The Properties of Gases and Liquids, 3rd edn., McGraw-Hill, New York, 1977. T. Booblik, V. Fried and E. Hala, Vapor pressures of pure substances, Elsevier, Amsterdam, 1973. H.C. Van Ness, S.M. Byer and R.E. Gibbs, AIChemE J., 19 (1973) 238-244. A. Fredenslund, J. Gmehling and P. Rasmussen, Vapor-Liquid Equilibria Using UNIFAC, Elsevier, Amsterdam, 1977. G.G. Beneto, A. Carton and M.A. Uruena, J. Chem. Eng. Data, 39 (1994) 249-250.
EQUIUBIilA ELSEVIER
FluidPhaseEquilibria127 (1997) 181-190
Thermodynamic properties of water + normal alcohols and vapor-liquid equilibria for binary systems of methanol or 2-propanol with water Brahim Khalfaoui a, *, Abdessalem H. Meniai b, Rafael B o r j a c a Chemical Engineering Department, UMIST, Sackville Street, Manchester, M60 1QD, UK b Institut de Chimie Industrielle, Universite de Constantine, 25000 Constantine, Algeria c Instuto de la Grasa y sus Derivados (CSIC), Avda. Padre Garcia Tejero 4, 41012 Seville, Spain
Received 15 January 1996; accepted 17 May 1996
Abstract
Experiments have been carried out to determine densities and refractive indices of binary and ternary mixtures of water with normal alcohols (methanol and n-butanol) at 293.15 K. Also, vapor-liquid equilibrium data have been measured for binary systems of methanol-water and 2-propanol-water at atmospheric pressure, from which the excess Gibbs free energy was computed for both systems. The experimental vapor-liquid equilibrium (VLE) data have been compared with those predicted by the Wilson, NRTL and UNIQUAC equations. Agreement has been found to be very satisfactory. Keywords: Methanol-H20 systems; 2-Propanol-H20 systems; Density; Refractive index; Gibbs free energy; Vapor-liquid equilibria; Water-n-alcohol systems
I. Introduction
Vapor-liquid equilibrium data are always required for engineering use such as in design and, especially, in distillation, which is the most common operation performed in the chemical industry for the separation of liquid mixtures. In previous work [1,2] we presented experimental results of vapor-liquid equilibrium (VLE) measurements for binary mixtures of water with chlorinated hydrocarbons and water with brominated hydrocarbons. The present work reports the VLE results of measurements on binary mixtures of methanol-water and 2-propanol-water at atmospheric pressure. The experimental results have been compared with those predicted by the Wilson [3], NRTL
* Corresponding author (on leave from Universite de Constantine, 25000 Constantine, Algeria). 0378-3812/97/$17.00 Copyright© 1997ElsevierScienceB.V. All rightsreserved. PII S0378-3812(96)031 29-9
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190
182
(non-random, two-liquid) [4] and UNIQUAC [5] equations. The methanol-water system has been investigated previously [6-9] and this system is, therefore, suitable for testing the performance of the VLE apparatus used in this work. However, new reliable data were obtained for this system. Additionally, new accurate and reliable VLE data for the 2-propanol-water system are also obtained in the present study and are compared with those of Wilson and Simon [10], Choffe et al. [11], Verhoeye [12], and Verhoeye and De Schepper [9]. Prior to the VLE measurements, the densities and refractive indices of methanol-water, n-butanol-water, methanol-n-butanol and of the ternary system methanol-n-butanol-water were measured at 293.15 K. This was done as a continuation of a long-term study on the thermodynamic properties of liquid mixtures containing water and normal alcohols in general. Those measurements of particular interest have been made in order to expand the range of mixtures for which the physical properties can be calculated. Furthermore, VLE measurements on binary and ternary mixtures composed of methanol, n-butanol and water were previously carried out [ 13]. The VLE results of this work were used to compute the excess Gibbs free energy for the two studied systems.
2. Experimental section 2.1. Materials Methanol and n-butanol were supplied by Fisons Scientific Equipment Ltd., and 2-propanol was supplied by BDH Chemicals Ltd.. The purity of these alcohols was checked by gas chromatography (GC) and was found to be in excess of 99%. The alcohols were used without further purification. The water used was doubly distilled tap water. Table 1 compares some of the measured properties with literature values. As can be seen, the agreement is good.
2.2. Apparatus and procedure The densities of the pure liquids and the mixtures were measured at 293.15 K using an Anton Paar DMA 46 densimeter with a precision of _0.0001 g cm -3. The refractive index measurements at 293.15 K were carried out using a dipping refractometer which was equipped with thermoprisms (Carl Zeiss, Jena) that had a measurement precision of + 2 × 10 -5. Illumination was provided by a sodium lamp. Both instruments were statically controlled at 293.15 + 0.05 K by circulating water from a
Table 1 Densities p, refractive indices r/ and normal boiling points Tb of the pure materials used in this study at 293.15 K Component p (g cm -3) 7/(D) Tb (K) Methanol 2-Propanol n-Butanol Water
Exptl. 0.7915 0.7853 0.8100 0.9980
Lit. 0.7915 [18] 0.78534 [ 16] 0.8098 [ 17] 0.9981 [20]
References are given in square brackets.
Exptl. 1.32900 1.37780 1.39913 1.33301
Lit. 1.32885 [19] 1.37745 [ 16] 1.39920 [ 17] 1.33320 [20]
Expti. 337.80 355.50 390.80 373.13
Lit. 337.68 [ 18] 355.41 [18] 390.15 [15] 373.20 [21]
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
183
thermostat. The VLE measurements were carried out using a vapor-liquid equilibrium apparatus. This consisted of an all-glass dynamic recirculating still described by Walas [ 14], equipped with a Cotterell pump. The apparatus used (Labodest model, manufactured by Fischer Labor und Verfahrenstechnic, Germany) is designed for a charge of approximately 100 cm 3 of mixture and can be operated at low, moderate or high pressures, ranging from 0.25 kPa through atmospheric pressure up to 400 kPa, and at temperatures up to 523.15 K. The vapor and liquid phases reach equilibrium once the vapor temperature has remained constant for a period of 15-30 min. At this moment, the first sample take-off can begin. The equilibrium temperature was measured with a digital Fischer thermometer with a precision of 0.01 K, and the pressure with a digital manometer with a precision of 0.01 kPa. For further details, the apparatus used in this work was described elsewhere by Aucejo et al. [15]. 2.3. Analysis The liquid and the condensed vapor compositions were determined using a Hewlett Packard 5890 S II gas chromatograph equipped with a thermal conductivity detector (TCD). The GC response peaks were integrated by using an electronic integrator. The packed chromatographic column used was a fused silica capillary column 10 m in length and 0.35 mm in internal diameter, coated with Poraplot Q. Previous calibration analysis was carried out with mixtures of known composition. Usually, a single analysis of the vapor or liquid composition by GC is imprecise. However, with repeated measurements, the standard deviation of a composition analysis was generally less than 0.001 mole fraction. At least two analyses were performed for each liquid and each vapor composition except for the methanol-water system, whose compositions were determined from both GC and density measurements in order to check the reproducibility of the analytical results.
3. Results and discussion
The physical properties of the pure components used in this work are given in Table 1 along with the literature values taken from Refs. [15-21]. The experimental densities and refractive indices of the three binary systems and the ternary system at 293.15 K are given in Table 2. The densities were measured with an uncertainty of _ 0.0002 g cm 3. The experimental VLE results for the two binaries, methanol-water and 2-propanol-water at atmospheric pressure together with the calculated values for the liquid phase activity coefficients (3,1 and 3'2) and the excess Gibbs free energy are given in Tables 3 and 4, respectively. The activity coefficients, 3'i, in the liquid phase were calculated using the following equation 3"i = Y i P / x i P°
(1)
where the vapor pressures pO were calculated with the Antoine equation given in Table 5 along with the Antoine constants taken from Ref. [22] by assuming ideal behavior of the mixture in the vapor phase. The Yi values computed using Eq. (1) are listed in Tables 3 and 4. It can be seen that both systems present a positive deviation from ideality. A comparison of the VLE results for the methanol-water system with the data cited earlier in this paper is made in Fig. 1, where the vapor phase composition is plotted against the liquid phase composition. As can be seen, the agreement is generally good considering the difficulties inherent in measuring vapor-liquid equi!ibria in the dilute
B. Khalfaoui et al./Fluid Phase Equilibria 127 (1997) 181-190
184
Table 2 Densities and refractive indices of solutions of water (1), methanol (2) and n-butanol (3) at 293.15 K W1
100 0.0 0.0 6.37 12.43 18.30 29.91 64.65 7.15 12.62 32.69 46.08 79.59 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 6.50 8.30 8.60 8.90 9.00 11.20 13.03 19.23 20.64 25.60 26.73 29.44 36.53 36.71 52.45 60.66 71.41
W2
0.0 100 0.0 93.63 87.57 81.70 70.09 35.35 0.0 0.0 0.0 0.0 0.0 92.15 74.50 60.19 14.35 12.75 7.45 2.75 1.43 81.20 17.10 35.20 52.30 84.90 67.60 32.57 49.52 38.83 33.22 53.07 22.65 52.44 41.07 34.97 21.49 6.56
W3
0.0 0.0 100 0.0 0.0 0.0 0.0 0.0 92.85 87.38 67.31 53.92 20.41 7.85 25.50 39.81 85.65 87.25 92.55 97.25 98.57 12.30 74.60 56.20 38.80 6.10 21.20 54.40 31.25 40.53 41.18 20.20 47.91 11.03 22.22 12.58 17.85 22.03
P
(g cm- 3)
(D)
0.9980 0.7915 0.8100 0.8201 0.8328 0.8430 0.86326 0.9350 0.8307 0.8417 0.8787 0.9066 0.9848 0.7971 0.7981 0.7990 0.8074 0.8076 0.8086 0.8096 0.8099 0.8111 0.8224 0.8197 0.8170 0.8112 0.8213 0.8347 0.8404 0.8492 0.8610 0.8659 0.8679 0.8876 0.8850 0.9181 0.9319 0.9665
1.3330 1.3290 1.3991 1.3298 1.3300 1.33 I0 1.3314 1.3320 1.3941 1.3911 1.3606 1.3501 1.3416 1.3405 1.3486 1.3568 1.3890 1.3902 1.3942 1.3982 1.3986 1.3401 1.3822 1.3710 1.3578 1.3336 1.3476 1.3725 1.3538 1.3659 1.3681 1.3490 1.3720 1.3490 1.3568 1.3563 1.3540 1.3438
r e g i o n s o f m i x t u r e s . T h i s w a s d o n e to e v a l u a t e the p e r f o r m a n c e o f t h e V L E a p p a r a t u s u s e d in this w o r k . T h e V L E r e s u l t s o f t h e 2 - p r o p a n o l - w a t e r s y s t e m , set o u t in T a b l e 4, a r e a l s o p r e s e n t e d in g r a p h i c a l f o r m as a p l o t o f v a p o r p h a s e c o m p o s i t i o n v e r s u s l i q u i d p h a s e c o m p o s i t i o n in F i g . 2, w h e r e the a z e o t r o p i c p o i n t is s i t u a t e d at x = y = 0.689. A s c a n a l s o b e o b s e r v e d , the V L E r e s u l t s o b t a i n e d in this s t u d y a n d s h o w n o n the p l o t a r e in g o o d a g r e e m e n t w i t h t h o s e f o u n d in the l i t e r a t u r e c i t e d in the
185
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
Table 3 Vapor-liquid equilibrium data, pressure P, temperature T, liquid phase mole fraction x~, vapor phase mole fraction Yl, activity coefficients Yi and excess Gibbs free energy G E / R T , for the methanol (1) + water (2) system P (kPa)
T (K)
x1
Yl
3Jl
T2
GE/RT
99.988 99.988 100.042 99.782 100.135 100.155 100.135 100.228 100.682 100.727 100.737 100.737
337.80 337.85 337.90 337.95 340.85 344.65 347.55 350.65 354.55 357.75 362.65 373.15
1 0.8948 0.8150 0.7649 0.6060 0.4709 0.3477 0.2561 0.1757 0.1174 0.0574 0.0
1 0.9498 0.9115 0.9005 0.8309 0.7766 0.7078 0.6472 0.5852 0.4928 0.3742 0.0
1.0428 1.0972 1.1496 1.2002 1.2491 1.3835 1.5348 1.7672 1.9920 2.6174 -
1.9386 1.9401 1.7081 1.5286 1.2756 1.1964 1.11 44 1.0136 1.0196 0.9742 -
0.0 0.1071 0.1982 0.2325 0.2777 0.2335 0.2300 0.1903 0.1112 0.098 0.0306 0.0
Table 4 Vapor-liquid equilibrium data, pressure P, temperature T, liquid phase mole fraction Xl, vapor mole fraction y~, activity coefficients ~/~ and excess Gibbs free energy G E / R T for the 2-propanol (1) + water (2) system P (kPa)
T (K)
xI
Yl
Yl
72
GE/RT
102.27 102.27 102.27 102.12 102.27 102.25 102.13 102.40 102.03 101.95 101.98 101.98
355.5 352.35 352.85 353.15 353.25 353.55 354.05 354.45 356.35 356.95 357.75 373.13
1 0.8495 0.7975 0.7541 0.6628 0.4887 0.3241 0.2407 0.1350 0.0840 0.0456 0.0
1 0.8015 0.7512 0.7201 0.6854 0.6150 0.5641 0.5517 0.4764 0.4582 0.4380 0.0
1.0799 1.0565 1.0565 1.1412 1.3717 1.8575 2.4128 3.4309 5.1743 8.8313 -
2.945 2.6883 2.4568 2.0086 1.6011 1.3421 1.2123 1.1 48 1 1.0950 1.0560 -
0.0 0.2278 0.2440 0.2624 0.3227 0.3951 0.3995 0.3582 0.2859 0.2211 0.1513 0.0
Table 5 Antoine constants of materials used in VLE
a
Component
A
B
C
Methanol 2-Propanol Water
8.07246 7.56634 7.96681
1574.990 1366.142 1668.210
238.860 197.970 228.00
a logloPiO= A - B / ( t + C ) , Centigrade.
where the units of Pi0 are millimeters of mercury and the temperature is in degrees
186
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190
e II:F 06
'~
y
05
,
/
!
i
C4
p
L
//
02
/
/
/
)/
C.O 0.0
]
i 02
0 4
08
06
10
X 1
Fig. 1. Vapor-liquid equilibrium compositions for methanol-water at atmospheric pressure. (O) This work; ([]) Hala et al. [6]; (zx) Dalager [7]; ( 0 ) Maripuri and Ratcliff [8]; ( v ) Verhoeye and Schepper [9].
introduction. Because of this good agreement, it can be concluded that the experimental errors are trivial for both systems. The excess Gibbs free energy was obtained using the following thermodynamic relationship (2)
G S = RT~_+xi ln yi i
I.C
7
r
-
-
T
. . . . .
~
-
-
08
oo
+'+ /
//
/ 0.0 0.0
0.2
I
I
[
0,4
O.6
0.8
X 1
Fig. 2. Vapor-liquid equilibrium compositions for 2-propanol-water at atmospheric pressure. (O) This work; ([q) Wilson and Simon [ 10]; (zx) Choffe et al. [11 ]; (O) Verhoeye [12]; ( v ) Verhoeye and Schepper [9].
B. Khalfaoui et al. /Fluid Phase Equilibria 127 (1997) 181-190
187
0.3
0.2 E--
0.1
i
0.0( 0.0
0.2
I 0.4 X1
I
I
0.6
0.8
1.0
Fig. 3. Excess Gibbs free energy as a function of the liquid mole fraction for the methanol (1) + water (2) system. (O) Experimental points; ( - - ) splined curve.
and the results for 2-propanol-water and m e t h a n o l - w a t e r are shown in Figs. 3 and 4, respectively. As can be seen from these figures, the two systems exhibit a behavior that is slightly unsymmetrical. The excess Gibbs free energy expressed as GE/RT shows a maximum almost at x I = x 2 for the two systems. The fitness of the curves is an indication o f the precision of the data obtained; therefore, it can be seen in Figs. 3 and 4 that good results have been achieved for the two systems. The results have been tested for thermodynamic consistency using the point-to-point method o f Van Ness et al. [23] modified by Fredenslund et al. [24]. A four-parameter Legendre polynomial was
,
i
1
i
0.4
c.3
L~
0
0.2
0.1
0.0
0.0
t
i
J
,
0.2
0.4
O.6
0.8
1.0
x 1
Fig. 4. E x c e s s G i b b s f r e e e n e r g y as a f u n c t i o n o f the l i q u i d m o l e f r a c t i o n f o r the 2 - p r o p a n o l (1) + w a t e r ( 2 ) s y s t e m . (o) E x p e r i m e n t a l points; ( - - ) splined curve.
B. Khalfaoui et a l . / Fluid Phase Equilibria 127 (1997) 181-190
188
Table 6 Parameters of the Wilson, NRTL and UNIQUAC models
Model Wilson NRTL UNIQUAC
Parameters
Methanol (1) + water (2)
(cal mo1-1 )
NRTL (al2 = 0.30)
MAD(y)
NRTL (oil2 ~ 0.45)
MAD(y)
3,12- 3,22 3,21- Al i g 12- g22 g21- gJl
- 714.80 1805.11 334.44 326.99 1462.45 - 539.00
0.010
756.23 1292.24 701.00 1726.71 562.51 - 82.86
0.010
ul2-u22 u21- u ll
2-Propanol (1) + water(2)
0.009 0.008
0.009 0.009
used for the excess G i b b s free energy. According to Fredenslund et al., the P - T - x - y data are consistent if the m e a n absolute deviation between calculated and measured m o l e fractions o f c o m p o n e n t 1 in the v a p o r p h a s e ( M A D ( y ) ) is less than or equal to 0.01. The values o f M A D ( y ) , for both systems, were found to lie within the stated tolerance, as can be seen in T a b l e 6. As indicated by this test, it can be concluded that our results are consistent. T h e activity coefficients for the two binaries investigated in this study, given in Tables 3 and 4, are c o m p a r e d with those predicted b y the Wilson, N R T L and U N I Q U A C equations and are shown in Figs. 5 and 6 as plots o f In 'Yi versus the liquid phase composition. As can be seen f r o m both graphs, the a g r e e m e n t is satisfactory for the three models. Moreover, as can be observed, the change in In 'Yi with c o m p o s i t i o n is an adequate demonstration o f the t h e r m o d y n a m i c consistency o f the obtained results. The adjustable binary
1.2
1 3
I
0.8 L-'N~X 2 ~.~. ,n71
InT2
I
3-~
/y
x1
Fig. 5. Activity coefficients as a function of liquid compositions for the methanol (1) + water (2) system. (O, O) This work. Curve 1, Wilson; curve 2, NRTL; curve 3, UNIQUAC.
B. Khalfaoui et al. / Fluid Phase Equilibria 127 (1997) 181-190
25~_
2~
2.0 I_~o
t
,
~
1 5
00
189
~
02
0 a xI
0.6
0 g
1.0
Fig. 6. Activity coefficients as a function of liquid compositions for the 2-propanol (1) + water (2) system. ( O , 0 ) This work. Curve 1, Wilson; curve 2, NRTL; curve 3, UNIQUAC.
parameters of each model were estimated in the same manner as performed by Benito et al. [25], by using a non-linear regression method that minimizes the following objective function N
E i= l
[() ")/1
"Yl
In - -
"Y2
-In i,calc
-")/2
(3) i,expt
The resulting parameters from the fit to each model are shown in Table 6.
4. List of symbols A,B,C gij
P
P? R t
T uij
W xi Yi
Antoine constants in Table 5 interaction parameter in NRTL equation given in Table 6 pressure pure component vapor pressure gas constant in Eq. (2) temperature (°C) absolute temperature interaction parameter in UNIQUAC equation given in Table 6 weight percentage in Table 2 liquid phase mole fraction vapor phase mole fraction
4.1. Greek letters Ol-ij
")'i
non-randomness parameter in NRTL equation given in Table 6 activity coefficient
19o
B. Khalfaoui et al./ Fluid Phase Equilibria 127 (1997) 181-190 refractive index interaction e n e r g y in W i l s o n e q u a t i o n g i v e n in T a b l e 6 density (g crn - 3 )
Aij P
4.2. Subscripts i, 1,2,3
c o m p o n e n t identification
4.3. S u p e r s c r i p t s o
pure c o m p o n e n t
References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25]
M.R. Cooling, B. Khalfaoui and D.M.T. Newsham, Fluid Phase Equilibria, 81 (1992) 217-229. B. Khalfaoui and D.M.T. Newsham, Fluid Phase Equilibria, 98 (1994) 213-223. G.M. Wilson, J. Am. Chem. Sot., 86 (1964) 127-130. H. Renon and J.M. Prausnitz, AIChE J., 14 (1968) 135-144. D.S. Abrams and J.M. Prausnitz, AIChE J., 21(1) (1975) 116-128. E. Hala, J. Wichterle, J. Polak and T. Boublik, Vapour-Liquid Equilibrium Data at Normal Pressures, Pergamon Press, London, 1968. P. Dalager, J. Chem. Eng. Data, 14 (1969) 298-301. V.O. Maripuri and G.A. Ratcliff, J. Chem. Eng. Data, 17 (1972) 366-369. L. Verhoeye and H.De Schepper, J. Appl. Chem. Biotechnol., 23 (1973) 607-619. A. Wilson and E.L. Simon, Ind. Eng. Chem., 44(2) (1952) 2214-2219. B. Choffe, M. Cliquet and S. Meunier, Rev. Inst. Ft. Pet., 15(6) (1960) 1051-1060. L.A.J. Verhoeye, J. Chem. Eng. Data, 13 (1968) 462-467. D.M.T. Newsham and N. Vahdat, Chem. Eng. J., 13 (1977) 27-31. S.M. Walas, Phase Equilibria in Chemical Engineering, Butterworth, London, 1985. A. Aucejo, M.C. Burguet, J.B. Monton, R. Munoz, M. Sanchotello and M.I. Vazquez, J. Chem. Eng. Data, 39 (1994) 271-274. A. Aucejo, V.G. Alfaro, J.B. Monton and M.I. Vazquez, J. Chem. Eng. Data, 40 (1995) 332-335. A. Weissberger, E.S. Proskauer, J.A. Riddick and E.E. Toops, Jr., Organic Solvent, Interscience, London, 1955. J. Timmermans, Physico-Chemical Constants of Pure Organic Compounds, Vol. 1. Elsevier, New York, 1950. R.C. Weast (Ed)., Handbook of Chemistry and Physics, The Chemical Rubber Co., Cleveland, OH, 1968. E.G. Mahers and R.A. Dawe, J. Chem. Eng. Data, 31 (1986) 28-31. R.C. Reid, J.M. Prausnitz and T.K. Sherwood, The Properties of Gases and Liquids, 3rd edn., McGraw-Hill, New York, 1977. T. Booblik, V. Fried and E. Hala, Vapor pressures of pure substances, Elsevier, Amsterdam, 1973. H.C. Van Ness, S.M. Byer and R.E. Gibbs, AIChemE J., 19 (1973) 238-244. A. Fredenslund, J. Gmehling and P. Rasmussen, Vapor-Liquid Equilibria Using UNIFAC, Elsevier, Amsterdam, 1977. G.G. Beneto, A. Carton and M.A. Uruena, J. Chem. Eng. Data, 39 (1994) 249-250.