Vapor-Liquid Equilibria 2-PROPANOL- WATER SYSTEM ABRAHAM WILSON' AND EDWARD L. SIhl0NS2 Rutgers University, New Brunswick, N . J .
AS
2. The operating pressure was controlled with a stainless steel Cartesian manostat, P (Emil Greiner Model 5 ) , and was measured with a Merriam high pressure manometer (Model A-995WM in series with an Ashcroft dead-weight gage tester (catalog No. 1300-3), as described by Othmer and Silvis (16). 3. Although the still was originally charged through bomb M , it was later found more satisfactory to introduce the charge into the previously evacuated still through valve I. 4. After temperature and pressure equilibrium had been maintained in the still for an hour, the distillation was stopped suddenly by admitting nitrogen to a pressure about 10 pounds per square inch greater than the operating pressure and shutting off the heaters. Samples of pot liquid and vapor condensate were then withdrawn into evacuated 50-cc. stainless steel sample bombs attached to valves I and 11. The introduction of excem nitrogen pressure prevented flashing of the still contents upon removal of the liquid samples. However, the liquid samples so obtained were saturated with nitrogen a t 4 to 5 atmospheres pressure. In order to remove the dissolved gas without IOSP of sample the Eample bomb was attached to a previously evacuated 600-cc. bomb, the connecting valve opened, and thc entire a@sembly placed in a refrigerator for an hour. The nitrogen pressure in the assembly was thus reduced to less than 1 atmosphere, and after admitting dry air to the bomb the liquid sample could be removed without change in composition. Tests with solution. of known composition, saturated with nitrogen a t 5 atmoqpheres, confirmed the effectiveness of this technique.
PART of a program t o study the effect of pressure upon the behavior of binary azeotropic systems, a study of the vapor-liquid equilibrium relationships in the system Zpropanolwater has been carried out over a range of pressures from 95 mm. t o 4 atmospheres. The only previously reported investigations of this system are at 760 mm. by Schumaker and Hunt ( $ I ) , Langdon and Keyes (8),Lebo ( I O ) , and Brunjes and Bogart ( 4 ) , and of these only the latter two include liquid boiling points. The data herein reported for 760 mm. are in good agreement with those of Brunjes and Bogart. MATERIALS
The water was distilled from a Barnstead laboratory still and stored in a 20-liter borosilicate glass carboy. C.P. grade 2-propanol was dried over calcium oxide and twice distilled through a 30-plate Brunn column a t a reflux ratio of 10: 1. A middle fraction, boiling in the range 82.23" to 82.28" C. a t 760 mm., was collected and stored in brown glass bottles. Bmnel, Crenshaw, and Tobin (3) report a mean boiling point for a number of determinations of 82.28' C. and aprobable boiling point of 82.26" C. The purified alcohol had a density of 0.78091 gram per ml., compared with Brunel's value of 0.78084 gram per ml. for pure 2propanol. Considering the principal impurity to be water, the mole per cent of the 2-prppanol was calculated as 99.98 on the basis of the change in density with composition a t 100.00% reported by Lebo (10). The compositions of the synthetic mixtures prepared for analytical purposes were calculated on the basis of this figure for the alcohol.
A detailed description of the complete apparatus and its opersr tion, including a complete legend for Figure 1, has been reported
TABLEI. SPECIFIC GRAVITIESO F LTIXTURES AND
ANALYTICAL METHOD
Equilibrium samples of liquid and vapor condensate were malyBed by specific gravity measurements made using a standingtype precision 20-ml. pycnometer of borosilicate glass (Eck and Krebs, catalog No. 4154). Analytical checks using this apparatus never exceeded and were usually less than 1 0 . 2 mg. Values of the specific gravities of known mixtures are presented in Table I. The constant temperature bath used in these measurements was controlled t o rt0.01 O C. by a mercury-toluene regulator.
TABLE 11.
OF %PROPANOL
WATER
2-Propanol, blole yo
d z 6 , G./hlI.
0.00 10.10 20.11 29.93 40.15 60.01 59.98 70.11 80.12 89.17 99,98
0.98708 0.95356 0,91254 0.88166 0.85696 0.83847 0.82314 0,81011 0.79912 0.78O78 0,78091
I r A P O R PRESSURES O F 2-PROPANOL I N THE 1 TO 8 ATJIOSPHERES
EXPERIMENTAL METHODS
RANGEO F
p , hlm.
Tlie subatmospheric measurements n-ere made as described by 13:ichman and Simons (1). The superatmospheric determination8 uiare carried out in an Othmer-hIorley high pressure still (16), niodified as shown in Figure 1 The principal changes were: 1. As suggested by Gilmont ( 6 ) , ballast tank E was disconuwted from trap D, and the nitrogen pressure was applied through the stainless steel reflux condenfier, R. After venting any nitrogen originally present in the still through valve IV, no additional gas was permitted to enter the system. Copper-constantan thermocouples a t positions 5 and 6 served as guides in keeping the nitrogen-vapor interface within the 20-inch long condenser.
Boiling Point,
c.
Exptl.
Calcd. from Eq. 1
81.61 84.85 86.56 89.07 91.79 93,50
742,4 843.2 916.8 994.9 1102.4 1174.4 1250.9 1357.9 1430.9 1507.9
742.4 844. 9 918.0 995.4 1102.4 1174.1 1260.7 1387.2 1431.0 1507.9
0
95.24
97.53 99,04 100,55
p , Lb./Bq. In. Abs.
Exptl.
1 Present address, Research and Development Division, Colgate-PalmolivePeet Co., Jersey City, N. J. 2 Present address, General Electric Research Laboratory, Schenectadg, N. Y .
2214
100.94 113.33 122.71 130.46 136.93 142.63 147.63
29,719 44.698 59.698 74.860 85,860 104,86 119.86
Calcd. f r o m Eq. 2 29.719 44,663 59.640 74.875 89.867 104.98 119.86
September 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY
elsewhere (93). Detailed operating instructions are available on request from one of the authors. Calculations of activity VAPORPRESSURE MEASUREMENTS. coefficients require vapor pressure data for the pure components. For water, the vapor pressure data given by Lange (9) were used. For 2-propanol, in the region from 35" t o 90" C., the data of Parks and Barton (18) were used. Above atmospheric pressure, the boiling points of 2-propanol were determined in the high pressure equilibrium still, and the data are given in Table 11. They can be satisfactorily represented by the following equations: ( 1 to 2 atmospheres) loglo p (mm.) =
-
+
(2 to 8 atmospheres) log,, p (pounds per square inch absolute) =
-
1593.60 [temp. ( " C.) 227.01
+
+ 6.33245
in which y = mole fraction in the vapor phase, 2 = mole fraction in the liquid phase, po = vapor pressure of pure component at the boiling point of the solution, and P = total still pressure. Subscript 1 refers t o the more volatile component, %propanol, while subscript 2 refers t o the less volatile component, water. Also included in Table I11 are the values of the multiplication factor, zi, proposed by Benedict, Johnson, Solomon, and Rubin ( 2 ) as a correction t o the activity coefficient for the deviation from ideality of the gaseous phase and the influence of pressure on on the liquid phase. It is defined as
+
846.79 [temp. ( " C.) 137.411 6.7392 (1)
(2)
VAPOR-LIQUID EQUILIBRIUM DATA
The experimental results obtained for the various isobars are presented in Table I11 along with the valucs of the activity coefficients, y D , calculated from the relationships:
2215
(5)
in which, for each component, +iis the moIar volume of the pure liquid at the solution boiling point and pi is the second virial coefficient of the vapor. I n the absence of complete data for the above terms, the z values were calculated by means of Scheibel's nomograph (90)from the vapor pressure and critical data of the pure components. Calculations showed that only for the &atmob phere isobar did the z term significantly affect the relationships subsequently derived between yr ( = y y z i ) and composition. The data for the 760-mm. isobar are plotted in Figures 2 and 3 as boiling point-composition and distribution diagrams. The corresponding curves for the other isobars are of the same form displaced only slightly from the 760-mm. curves. I n Figure 3 the circles represent the experimental data, while the curve is that calculated from smoothed values of the activity coefficients. Large scale plots of the calculated distribution curves and the boiling point curves were used to establish the values of the azeotropic compositions and boiling points given in Table IV.
Figure 1. High Pressure Still
INDUSTRIAL AND ENGINEERING CHEMISTRY
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INDUSTRIAL AND ENGINEERING CHEMISTRY
September 1952 I
I
I
I
I
1.0
I
I
I
I
I
2217
I
I
I
I
I
'..O o LIQUID 0
4
k-
801
I
0.0
VAPOR
. r \ &
QI
Figure 2.
TABLE IV.
I
0.2
1
4 0.9
1.0
Boiling Point-Composition Diagram at 760 Mm.
AZEOTROPIC COMPOSITIONS AND BOILINQPOINTS
Mm. 3087 760 380 190 95
P,
-
i
0.3 0.4 0.5 06 0.7 0.8 MOLE FRACTION 2-PROPANOL
Boiling Point, 0
c.
w0.3 0.4 0.5 0.6 0.7 0.8 0.9 MOLE FRACTION 2-PROPANOL IN LIQUID
0.1
0.0.0
2-Propano1, Mole yo 69.50 68.70 67.50 67.05 66.70
120.45 80.10 63.90 49.33 36.00
0.2
Distribution Diagram at 760 Mm.
Figure 3.
0
.O
Experimental data Calculated from smoothed activity coefficient values
0
THERMODYNAMIC CONSISTENCY O F THE DATA
2-PROPANOL
WATER
The internal consistency of the data was tested by the methods of Carlson and Colburn (6) and of Redlich and Kister (19). Both are based upon the Gibbs-Duhem equation which, for a n j homogeneous system at constant temperature and pressure, defines the relationship between the composition of the system and the values of any partial molar quantities of the components. Carlson and Colburn have utilized this equation t o establish a relationship between the constants in the van Laar equations which may be used t o represent the variation with composition of the logarithm of the activity coefficient of each component. The k a l forms of the van Laar equations may be represented as
01
"80
0.1
OZ
03
0.5
04
OS
0.7
08
OS
10
MOLE FRACTIOH 2-PROPAHOL 1H LIQUID
Figure 4. Carlson and Colburn Test of Thermodynamic Consistency Experimental d a t a 760 - Plot of van Laar Equations 6 and 7
0
at
mm.
and that the variation with composition of log yl/y2 may be represented by the equation
B (7)
log YI/W
B(1 - 221) f C[621(1 - 21) - 11
i=.
D(l I n Figure 4 the variation of log y with composition is shown, the circles representing the data for the 760-mm. isobar, the solid lines the van Laar equations. A similar fit between the data and the equations was found for all the other isobars. The ability t o represent, with the same set of constants, the change in log yt and log 7 2 with composition is an adequate demonstration of the thermodynamic consistency of the data. The terminal values of A and B, which are listed for all the isobars in Table V, undergo a systematic change with pressure, the values of A increasing and those of B decreasing with increasing pressure. A similar. trend is noted in the acetone-water system (14. Redlich and Kister have shown t h a t the criterion for thermodynamic consistency may be expressed in the form
+
- 2 ~ 1 ) [ 1- 821(1 -
21)1+.
..
(9)
in which B, G, and D are adjustable constants. The variation of log y1/y2 with composition is shown in Figure 5 in which the circles represent the 760-mm. data, and the solid line is a plot of Equation 9 adjusted t o satisfy Equation 8-i.e., the total area under the curve is zero. The data for all the isobars could be satisfactorily represented by Equation 9, although for 4 atmos-
TABLE V.
CONSTANTS IN THE VAN LAARA N D THE REDLICH AND KISTEREQUATIONS Redlich and Kister
van Laar
P, Mm. 3087 760 380 190 95
A 1.020 1.000 0.966 0.930 0.918
B 0.446 0.483 0.520 0.532 0.546
B 0.625 0.647 0.664 0.665 0.666
C -0,240 -0.206 -0.212 -0,200
-0.194
D 0.12Q 0,076 0.099 0,077 0.065
INDUSTRIAL AND ENGINEERING CHEMISTRY
2218
Vol. 44, No. 9
the change is too small t o be represented by any of the empirical relationships which have been proposed between azeotropic composition and pressure (7, 13, 17, 22). Licht and Denaler ( 1 2 ) have shown, on examination of 18 twocomponent azeotropic systems, that the boiling temperature of an azeotrope is a unique and single-valued function of the pressure, independent of changing composition. Over a limited range at least, the relationship can be expressed by
MOLE
Figure 5.
FRACTION PPROPANOL
Redlich and E s t e r Test of Thermodynamic Consistency 0 Experimental data at 760 mm.
- Plot of Equation 9
141
0 AZEOTROPE 0 2-PROPANOL
\\
-I
WATER
1.2 --
W u)
a 1.1 w
log Pr = A log Pa
+B
(11)
in which A and B are constants. The linear variation of log Pa: with log Pr is shown in Figure 7 in li-hich 2-propanol is the
13-
-
in which T is the absolute temperature and A and B are constants. Plots of the logs of the vapor pressure against the reciprocal of the absolute temperature for the azeotrope, water, and 2-propanol are shown in Figure 6. All three lines, although exhibiting a slight curvature, are practically parallel. Since a requirement for the disappearance of the azeotrope is t h a t its vapor pressure curve cross that for one of the components, it is unlikely t h a t suah a disappearance of the 2-propanol-water azeotrope will occur a t a n y reasonable pressure (13). If the azeotropic vapor pressure curve can be assumed t o parallel the experimental 2-propanol vapor pressure curve in the range of 2 to 8 atmospheres, the former may be readily extrapolated t o those higher pressures. Othmer (14) has suggested a n empirical relationship between the isothermal vapor pressures of the azeotrope, Pa, and of a reference substance, Pr,
-
i:
0.9-
4
0.8-
I
5 0.70
5 0.6-
03
2l
O0.I
o .d
023
I
I
026
027
I I 0.28 029 1/Tx102
I 030
I 031
I
0 3 2 '
Figure 6. Variation with Temperature of Vapor Pressures of Water, 2-Propanol, and Azeotrope
pheres and 760 mm. the calculated curves deviated from the experimental curves in the regions of high and lorn 21 values, the fit being good over the rest of the concentration range. Apparently a t the higher pressures the three-constant equation cannot be used t o represent adequately the experimcntal data. No attempts were made t o obtain a better fit with a more complex equation. The constants B, C, and D for each of t h e isobars are listed in Table V. EFFECT OF PRESSURE ON AZEOTROPIC BEHAVIOR
As seen from the data in Table IV the mole fraction of %propanol in the azeotrope and the azeotropic boiling point both increase with increasing pressure. The change in composition is of the same magnitude, though opposite in sign, as t h a t observed in the acetone-carbon tetrachloride system ( I ) , and as in t h a t case,
LOG ( V . P . x I 0 ATMOSJ AZEOTROPE
Figure 7. Variation of 2-Propanol Vapor Pressure with That of Azeotrope
reference liquid. This relationship, like the one mentioned previously, may be used t o extrapolate the effect of pressure on the boiling point of the azeotrope in the range in which vapor pressure data exist for the reference substance. When decadic logarithms are used, t h e values of A and B for this system are 1.0148 and -0.03914, respectively. ACKNOWLEDGMENT
The authors gratefully acknowledge the support given to this investigation by the Research Council of Rutgers University.
September 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY NOMENCLATURE
A , B, C, D = constants d = density, grams per ml. p = vapor pressure, mm. or pounds per square inch absolute po = vapor pressure of pure components at boiling point of solution P = total still ressure Pa = isothermarvapor pressure of azeotrope Pr = isothermal vapor pressure of reference substance R = ideal gas constant T = absolute temperature z = mole fraction in liquid phase y = mole fraction in vapor phase z, = multiplication factor & = molar volume of pure liquid at solution boiling point 6% = second virial coeflicient of vapor y o = activity coefficient yi = activity coefficient, corrected
Subscripts 1 refers to more volatile component, %propanol 2 refers t o less volatile component, water LITERATURE CITED
(1) Bachman, K. C., and Simons, E. L., IND. ENG.CHEM.,44, 202 (1952). (2)Benedict, M.,Johnson, C. A,, Solomon, E., and Rubin, L. C., Trans. Am. Inst. Chem.. Engrs., 41, 371 (1945).
2219
(3) Brunel, R. F., Crenshaw, J. L., and Tobin, E., J. Am. Chem. SOC.,43, 561 (1921). (4) Brunjes, A. S.,and Bogart, M. J. P., IND. ENG.CHEM..35, 255 (1 943). (5) Carlson, H. C., and Colburn, A. P., Zbid., 34, 581 (1942). (6) Gilmont, R.,private communication. (7) Horsley, L. H., Anal. Chem., 19,603 (1947). (8)Langdon, W.M., and Keyes, D. B., IND.ENG.CHEX.,34, 938 (1942). (9) Lange, N. A., “Handbook of Chemistry,” 6th ed.. pp. 1443, 1452, Sanduaky, Ohio, Handbook Publishers, Inc., 1946. (10) Lebo, R. B., J . Am. Chem. Soc., 43, 1005 (1921). (11) Lecat, M., 2. anorg. u. aZZgem. Chem., 186, 119 (1930). (12) Licht, W., Jr., and Denzler, C. G., Chem.. Eng. Progress, 44, 627 (1948). (13) Nutting, H. S.,and Horsley, L. H., Anal. Chem., 19,602 (1947) (14) Othmer, D. F., IND.ENQ.CHEM.,32, 841 (1940). (15) Othmer, D. F.,and Morley, F. R., Zbid., 38, 751 (1946). (16) Othmer, D. F.,and Silvis, S. J., private communication, (17) Othmer, D. F.,and Ten Eyck, E. H., Jr., IND.ENG.CHEM.,41, 2897 (1949). (18) Parks, G. S.,and Barton, B., J . Am. Chem. SOC., 50,24 (1928). (19) Redlich, O.,and Kister, A. T., IND.ENQ. CHEM.,40, 345 (1948). (20) Scheibel, E. G.,Zbid., 41, 1076 (1949). Zbid., 34,701 (1942). (21) Schumaker, J. E., and Hunt, H., (22) Skolnik, H., Ibid., 43, 172 (1951). (23) Wilson, A., thesis, Rutgers University, 1951. (24)Young, S.,and Fortley, F. C., #J. Chem. SOC.,1902, 728. RECEIVED for review December 20, 1851. ACCEPTEDMarch 26, 1952. Detailed operating instructions for the high pressure still are available upon request from Edward L. Simons.
Phase Behavior in the Hydrogen Sulfide-Water System F. T. SELLECK, L. T. CARMICHAEL, AND B. H. SAGE California Znstitute of Technology, Pasadena 4 , Calif.
B
O T H hydrogen sulfide and water are encountered in many petroleum reservoirs. A knowledge of the phase behavior of mixtures of these two components forms a part of the background of information requisite for an understanding of effective means of producing petroleum. The early investigations of the hydrogen sulfide-water system were directed primarily to a study of hydrates and t o the development of generalizations of phase behavior for binaty systems. D e Forcrand and coworkers (8, 3, 6) studied the hydrogen sulfide-water system in some detail, placing primary emphasis on the properties of the hydrates. Villard reviewed the history and made studies of this and other systems forming hydrates (22). Scheffer studied this system in some detail, determined the three-phase pressures with accuracy (16),and followed this experimental investigation with an overall review (17)at states involving solid phases. At a much later date Scheffer and Korvezee established with some certainty the composition of the hydrate (9)and confirmed the earlier finding of de Forcrand (4)t h a t the hydrate contains 6 molecules of water per per molecule of hydrogen sulfide. The higher values obtained earlier (9,6) appear to have resulted from the occlusion of water in the hydrate. Only limited data concerning the solubility of hydrogen sulfide in water have been obtained ( 7 ) . The quadruple point was well established (15-17) by Scheffer. Schreinemakers ($0)discussed the significance of the quadruple point and the triple point curves determined the presin this system, and Wright and Maass (8.4) sures and temperatures for equilibrium of the hydrate, the aqueous liquid, and the gas phase at temperatures near the freezing point of water.
The volumetric and phase behavior of water has been investigated in detail and was summarized by Keenan and Keyes (8). These data have been employed in the present study and no reference to the basic investigations is made. Hydrogen sulfide has been studied less extensively. Murphy (11) determined the volumetric behavior and measured the vapor pressure of this compound, and West ($3) summarized its thermodynamic properties. Additional data concerning the volumetric behavior and vapor pressure recently became available ( 1 2 ) . The present study relates t o measurements of the three-phase pressures and temperatures which are associated with t h e quadruple point found by Scheffer (15) at a temperature of 85.1’ F. and a pressure of 325 pounds per square inch. The composition of the coexisting phases in a two-phase equilibrium involving an aqueous liquid and a gas phase for temperatures between 100’ and 340” F. was determined. The measurements were made at pressures as high as 5200 pounds per square inch. The composition of the hydrogen sulfide-rich liquid was also measured a t temperatures below the critical temperature of t h a t compound. METHODS AND APPARATUS
The equipment employed was similar t o t h a t utilized in an earlier study (19) of nitrogen dioxide, with the exception t h a t a somewhat smaller spherical pressure vessel, for brevity called a bomb, was used in the present investigation. The details of the construction of the bomb are shown in Figure 1. An unsupported-area seal was provided t o close the threaded joint. The design of the seal is depicted in an enlarged insert in this figure. Lead sealing rings were employed in the present investigation, al-
