A. N. CAMPBELL AND G. M. MUSBALLY Department of Cl~etnistry, Utzioersity of Manitoba, Winnipeg, Manitoba Received May 20, 1970

Transcription

1 Vapor pressures and vapor * liquid equilibria in the systems: (I) acetone - chloroform, (2) acetone - carbon tetrachloride, (3) benzene - carbon tetrachloride A. N. CAMPBELL AND G. M. MUSBALLY Department of Cl~etnistry, Utzioersity of Manitoba, Winnipeg, Manitoba Received May 20, 1970 I I i The saturation vapor pressures of ten mixtures of the binary systems (I) acetone -chloroform, (2) acetone - carbon tetrachloride, and (3) benzene - carbon tetrachloride have been determined, from 100 to 230" for system 1 and from 100" up to the highest temperature at which llquid and vapor coexist for systems 2 and 3. The system acetone - chloroform could not be studied at higher temperatures because of decomposition. The gas-liquid critical temperatures of the three binary systems have been determined by the disappearance of meniscus method. The orthobaric compositions of the vapor-liquid equilibria of the binary systems have been measured from 100 to 180' for system 1 and from 100" to the critical region for systems 2 and 3, using a glass bomb enclosed in a steel bomb. From the vapor-liquid composition curves and the vapor pressure curves at constant temperatures (100, 150, 160, 170, and 180 ), the existence of an azeotrope in the system acetone-chloroform at these temperatures, and having a composition of 36.2 mole% acetone at low, was confirmed. The composition of the azeotrope shifts towards lower acetone content as the temperature is raised. Azeotropes were not found In the systems acetone - carbon tetrachloride and benzene - carbon tetrachloride, over the ranges of temperature and pressure of this research. The data of the binary systems were treated thermodynamically to yield the liquid phase activity coefficients and, as suggested by Chueh and Prausnitz, the Redlick-Kwong equation was used in a modified form to obtain the fugacity coefficients of components in the vapor phase. Several liquid phase parameters, such as the binary interaction constant, Henry's constant, and dilation constant have been calculated, using the van Laar equation as modified by Chueh and Prausnitz. I Canadian Journal of Chemistry, 48, 3173 (1970) Introduction Although most of the common liquids have been investigated at high temperatures, the thermodynamic treatment of mixtures has been largely limited to those of the paraffins. The reason for this is obviously the difficulties of the technique. Acetone, chloroform, benzene, and carbon tetrachloride were chosen for this study because they have been extensively investigated at low temperatures; the vapor P liquid equilibrium compositions have been determined for the binary systems between 10 and 55" and the vapor pressures at 25". Chatterjee and Campbell (1) have recently determined the orthobaric densities and vapor pressures of the pure liquids, up to the critical region. In addition to investigating the vapor pressures and vapor * liquid compositions, of the three binary systems, we have determined the critical temperatures of each system, as a function of composition. Rowlinson (2) points out that the critical temperature of a simple mixture is one of the most direct sources of information about the energy of interaction of two unlike molecules. he svstem acetone - chloroform has been thoroughly investigated at low temperatures. Campbell, Kartzmark, and Friesen (3) reported HE values for 25" and Campbell and Kartzmark (4) showed that a compound of acetone and chloroform exists in the liquid state; from their data on heats of mixing they calculated the enthalpy of the hydrogen bond. The vapor pressures and vapor * liquid compositions, at temperatures ranging from 15 to 55", have been studied by many workers, including Zawidski (5) and Rock and Schroder (6). A complete review of the work done on this system is given by Chatterjee and co-workers (7). The vapor P liquid compositions of the system benzene - carbon tetrachloride, at temperatures ranging from 25 to 45", have been investigated by at least 12 different investigators, including Young (8), Scatchard, Wood, and Mochel (9), and Fowler and Lim (10). The existence or nonexistence of an azeotrope has been a question for more than 50 years. Young (8) first suggested the possibility of an azeotrope of composition mole % carbon tetrachloride at 760 mm Hg pressure. Scatchard, Wood, and Mochel (9) investigated this system under isobaric and isothermal conditions at 760 mm and 40" but failed to find any azeotrope. Campbell and Dulmage

2 3174 CANADIAN JOURNAL OF CHEMISTRY. VOL. 48, 1970 (1 1) found that the vapor phase is never richer in benzene than the liquid phase, that is, there is no reversal of composition as there would be in passing through an azeotropic point. Vapor + liquid compositions of the system acetone - carbon tetrachloride have been measured at 0" by Gerrits (12), at 50" by Severns et al. (13), and by many other workers at relatively low temperatures. Brown and Smith (14) carried out vapor + liquid measurements at 45" and concluded that an azeotrope exists, under a pressure of mm at a mole fraction of acetone of Ellis and Rose (15) have measured the vapor pressures of this system over the temperati~re range 5-25". Experimental Purificatiot~ of Materials Organic impurities were removed from spectroanalyzed acetone by treating it with silver nitrate solution, followed by sodium hydroxide. After shaking, the solution was filtered and distilled from anhydro~ls calcium sulfate. The portion distilling between 56.1 and 56.2", under an atnlospheric pressure of 746 nun Hg, was collected. A fresh batch was prepared every week. Since it is notorious that chloroform decomposes when exposed to air and light, a fresh batch of this substance was prepared every 3 days. To remove ethanol, chloroform was shaken several times with concentrated sulfuric acid and then washed with water. The product was then dried with calciun~ chloride and distilled fro111 phosphorus pentoxide. The final distillate was stored in a closed cupboard, in brown bottles wrapped in black paper. Thiophene-free benzene was shaken with successive portions of concentrated sulfuric acid until a yellow color was no longer produced in the acid layer, then with water, dilute sodium hydroxide, and again with water. It was finally dried over phosphorus pentoxide and distilled from metallic sodium. Carbon tetrachloride was boiled with dilute potassium hydroxide solution, then shaken with alkaline potassium permanganate solution, and finally distilled from barium oxide. The refractive indices and densities (at 25") of our products were: Substance tldz5 d Acetone Chloroform 1, Benzene Carbon tetrachloride Method of Analysis of Mixtures Mixt~ires were analyzed refracton~etrically or by density, after preparation ofcalibration curves. Generally, the error of analysis was about 0.1 %. Detern~itmtion of Critical Tetrrperatrrre If a sealed tube containing liquid and vapor is heated uniformly, and if the overall density is not too different from the critical density, it is observed that the meniscus becomes flatter and fainter until the critical temperature is reached, when the meniscus disappears in the body of the tube. It had been observed by many workers, including ourselves, that it is not necessary that the overall density be exactly the critical density for these phenomena, including the occurrence of fluctuating striae, to be observed. There has been much discussion as to whether the temperature of disappearance of the meniscus is the true critical temperature, but, at all events, it is what we describe here as such. The method indicated above was the one chosen by us for the investigation of the systems (I) acetone - chloroform, (2) benzene - carbon tetrachloride, (3) acetone - carbon tetrachloride. The experimental tubes were 15 cm long and had an internal diameter of 2 mm. The internal volume of each tube was determined as a function of length. The pure liquid components were degassed by repeated freezing in liquid nitrogen and pumping off to a high vacuum line. Still under vacuum a suitable amount of liquid 1 was transferred to the capillary tube. The length of the tube occupied by liquid was measured with a cathetometer and the mass determined from the known density. The second liquid was then distilled into the capillary and the height again measured. Any change in volume on mixing was considered to be negligible for our purposes. The tube was then sealed while in liquid nitrogen. The tube with its contents of known composition was then placed in a thermostat at 200 O C. Temperature was controlled with a solid state proportional temperature controller. Up to a temperature of 225", Dow Corning Silicone Oil 550 was used as thermostat liquid. Above this temperature, the eutectic mixture of potassium, sodium, and lithiuln nitrates was used. To prevent corrosion in the latter case, the temperature sensor was a thermistor sheathed in stainless steel and the heating element was encased in a non-corrosive nickel-chromeiron alloy. The temperature of the thermostat was then re-adjusted at higher temperatures. As the critical temperature was approached, the temperature intervals were reduced. The meniscus was viewed through a telescope placed at a distance of 2 m. The temperature at which the meniscus disappeared was measured with a copper-constantan multiple junction thernlocouple. The bath was then allowed to cool slowly and the temperature of reappearance of the meniscus noted. The mean of these two temperatures was taken as the correct temperature. The reproducibility was 0.05 to 0.1". Measrrremetzt of Eq~~~liDriurrz Vapor-Liclrrid Co1?7positiotrs This has been described previously (1). Meas~rretnet~t of Vapor Pressrrre The method has been described previously (1). Not more than 30 lnin could be allowed for the establishment of equilibrium, because chemical decomposition became appreciable with time. To avoid this incipient decomposition as far as possible, two sets of experiments were carried out. In one, the temperature was raised from 100 to 190" and then a fresh filling was used for the range 200 to 280". Because the volume of

3 CAMPBELL AND MUSBALLY: VAPOR PRESSURE AND VAPOR-LIQUID EQUILIBRIA FIG. 1. MOLE % CHCL, Gas s liquid critical temperatures of the acetone-chloroform system as a function of mole% chloroform. MOLE % ACETONE FIG. 2. Gas 8 liquid critical temperatures of the acetone carbon tetrachloride system as a function of mole% acetone. the vapor phase was always much smaller than that of the liquid, we thought ourselves justified in assuming that the composition of the liquid did not change appreciably with temperature. Results Critical Temperatures A. Pure Liquids Critical temperatures of pure liquids are listed in Table 1. B. Binary Systems Acetone-Clzloroform The gas & liquid critical temperatures have TABLE 1 Critical temperatures of pure liquids Substance Critical temperature ('C) been determined over the whole concentration range. These data, together with most of the other extensive data of this paper, are to be found in the Depository of Unpublished Dataof the National Research Council of Canada.' The results are expressed graphically in Fig. 1, which represents critical temperature us. mole % CHCI,. Chloroform and mixtures having high chloroform content develop a yellowish color on prolonged heating. For this reason, tubes containing such mixtures were not introduced into the bath until the temperature was only a few degrees below the critical point. Acetone - Carbon Tetrachloride The data are expressed graphically in Fig. 2. Acetone 235.OkO. 1 Carbon tetrachloride 283.2k0.1 'Photocopies may be obtained free of charge, upon Benzene 288.0k0.1 request, from the Depository of Unpublished Data, Chloroform 263.2k0.1 National Science Library, National Research Council of Canada, Ottawa 2, Canada.

4 CANADIAN JOURNAL OF CHEMISTRY. VOL. 48, lb b MOLE % CCL* FIG. 3. Gas liquid critical temperatures of the benzene - carbon tetrachloride system as a function of mole% carbon tetrachloride. Benzene - Carbon Tetrachloride The data are expressed graphically in Fig. 3. Vapor + Liquid Equilibrium Compositions Acetone-Chloroform The liquid equilibrium compositions of this system have been determined at 100, 150, 160, and 180". The data are to be found in the Depository of Unpublished Data. A typical graph, that for 160, is given as Fig. 4. Acetone - Carbon Tetrachloride The data have been obtained for 100, 150,200, 250, and 270". A typical graph, that for 200, is shown as Fig. 5. MOLE % ACETONE IN LlOUlD FIG. 4. The vapor-liquid equilibrium composition curve of the system acetone-chloroform at 160 OC. MOLE % ACETONE IN LlOUlD FIG. 5. The vapor-liquid equilibrium composition curve of the system acetone-carbon tetrachloride at 200 "C. Benzene - Carbon Tetrachloride The data have been obtained for 100, 108, 150, 200, 250, and 270". The data for 150, which are typical, are plotted in Fig. 6. Vapor Pressures Acetone-Chloroform The saturation vapor pressures have been determined over the whole concentration range in the temperature interval 100 to 180". The figures, together with all other vapor pressures, are to be found in the Depository of Unpublished Data. The results for mixtures containing 7.50 and mole % acetone are expressed graph-

5 CAMPBELL AND MUSBALLY: VAPOR PRESSURE AND VAPOR-LIQUID EQUILIBRIA 3177 despite the fact that the compositions of the ten mixtures varied from 5.00 to mole% carbon tetrachloride. MOLE % CPRBON TETRACHLORIDE IN LlOUlD FIG. 6. The vapor-liquid equilibrium composition curve of the system benzene -carbon tetrachloride at 150 "C. ically in Fig. 7. All intermediate compositions lie between these two limits. Acetone - Carbon Tetrachloride Figure 8 shows the plots of vapor pressure us. temperature for different constant compositions, at the bubble-points of the system. Benzene - Carbon Tetrachloride Ten mixtures of different compositions were investigated between the temperatures of 100 and 270'. Figure 9 shows the relation between temperature and pressure for two of these mixtures: fhe pressureseare very similar for all mixtures, Discussion Thermodynamic Analysis Thermodynamic reduction and correlation of vapor + liquid equilibrium data are common at low pressures but, at high pressures, little attempt has been made to reduce such data with thermodynamically significant functions. In high pressure phase equilibria it is not possible to make the simplifying assumptions commonly made at low pressures; in such systems both phases, vapor and liquid, exhibit large deviations from ideal behavior. We have treated our experimental data by the method of Chueh and Prausnitz (16). Vapor pressure non-idealities are expressed as fugacity coefficients (+i) and the liquid phase nonidealities as activity coefficients (y,). The Redlich- Kwong equation, modified by the introduction of binary interaction constants, has been used for vapor phase properties. The van Laar equation, modified to allow for the rapid change of liquid molar volume which occurs in the critical region, has been used to represent the effect of composition on liquid phase properties. The parameters we have determined are : (I) Henry's constant, H,!, To. (2) The binary interaction constants a,,(,), defined by the equation ge/rt(x19, + ~292) = -a22(1) b2 I I I I I I I I I I TEMPERATURE 1.C I FIG. 7. Bubble point pressures of constant composition mixtures as function of temperature of the system acetone~hloroform.

6 CANADIAN JOURNAL OF CHEMISTRY. VOL. 48, 1970 TEMPERATURE I Cl FIG. 8. Bubble point pressures of critical composition mixtures as functions of temperature of the system acetone - carbon tetrachloride. The compositions are 0.0, 6.22, 15.01, 24.89, 34.11, 46.80, 52.01, 72.02, 58.20, 86.70, and 100 mole% acetone. FIG. 9. Bubble point pressures of mixtures of constant composition of the system benzene -carbon tetrachloride. The lower curve represents a mixture containing 5.00 mole% carbon tetrachloride, the upper a mixture containing mole% carbon tetrachloride, curves of the remaining eight mixtures lie in between. a,,, defined by the equation where ge = molar excess Gibbs energy, q,, q, = the effective molal volumes; Vcl, VG = the critical molal volumes. (3) The dilation constant. For a given binary system, the above parameters depend only on temperature, so that isothermal experimental data are required. From the large scale plots of P us. T, for fixed liquid compositions and of vapor s liquid equilibrium compositions, P-x-y data were read off for each isotherm. These data were then treated by the programs of Chueh and Prausnitz (16) and the analysis carried out, using an IBM computer. As all details are given in ref. 16, there is no need to repeat them here. Tl~ert?zodynamic Coizsisteizcy Test of the Vapor s Liquid Equilibriunz Data Vapor s liquid equilibrium data are said to be thermodynamically consistent when they satisfy the Gibbs-Duhem equation. When the data satisfy this equation, it is probable, but by no means certain, that they are correct. However, if they do not satisfy this equation, they are certainly incorrect. Thermodynamic consistency tests have been described by many authors; a good review is given in the monograph by van Ness (17). Chueh, Muirbrook, and Prausnitz (18) have described a consistency test which is an extension to isothermal high pressure data of the integral (area) test given by Redlich and Kister (19) and by Herington (20) for isothermal low pressure data. Again details are out of place here. Chueh

7 CAMPBELL AND MUSBALLY: VAPOR PRESSURE AND VAPOR-LIQUID EQUILIBRIA and Prausnitz (16) have replaced the data plotting procedure by a regression analysis. Chueh and Prausnitz use the criterion that if the average deviation of the activity coefficients y, and y, differ by more than 3 %, the data are inconsistent. Discussion of Results Chloroform-Acetone Figure 1, which is a plot of T,, the temperature of disappearance of meniscus, us. mole % chloroform, shows that the system chloroform-acetone deviates from ideality. This has been observed previously by many workers, including Swietoslawski and Kreglewski (21) and Chatterjee and co-workers (7). The deviation from ideality is also apparent in the vapor pressure plots. The negative deviation from Raoult's law has been explained by the fact that a 1 : 1 hydrogen bonded compound exists in this system, and Campbell and Kartzmark (4) have determined the enthalpy of the hydrogen bond as -2.7 kcal mole-'. We doubt, however, if this compound continues to exist under critical conditions. The existence of an azeotrope has been confirmed from the vapor + liquid equilibrium composition curves, e.g. Fig. 4. It was found that the composition of the azeotrope, a negative azeotrope in the terminology of Rowlinson (2), which is 36.2 mole% acetone at 100, shifts towards lower acetone content as the temperature is raised (Table 2). Rock and Schroder (6) came to the same conclusion from their study of this system at 10 to 55". This observation also agrees with the well known fact that an increase of temperature of a negative azeotrope decreases the mole fraction of the component whose vapor pressure increases more rapidly with temperature. This system was investigated only up to 180 "C, at which temperature the azeotrope still exists, but it is known (22) that the azeotrope does not exist up to the critical point. The experimental data for vapor + liquid composition and saturation vapor pressure ofthis binary system were treated to yield the fugacity coefficients, 4, and 4,, and the activity coefficients, y, and y,, of the two components, in the vapor and liquid phases, respectively. For this purpose we used the programs of Chueh and Prausnitz (16). Here again, most of the results are to be found in the Depository of Unpublished Data but a sample calculation (for 160") is given in Table 3. The second round of fitting has been TABLE 2 Dependence of azeotropic composition on temperature in the system CHCI,(l) - acetone(2) Temperature ("C) Azeotropic composition (mole fraction acetone) omitted since it was identical with the first round. From Table 3 it is seen that, though yi -> 1, as xi -> 1, the system acetone-chloroform is not ideal at the temperatures and pressures of this work. The activity coefficients have been plotted as functions of mole fraction acetone in the liquid phase, for 160" only, in Fig. 10. The curves for the other temperatures are similar. The deviations between the experimental values of the activity coefficients and the calculated values are rather large but this is discussed later. Carbon Tetrachloride (I) -Acetone (2) The plot of T,, the temperature of disappearance of the meniscus, us. mole % acetone (Fig. 2) for this system shows a slight deviation from ideality but the vapor pressure plots show neither maximum nor minimum. The vapor F? liquid equilibrium curves (e.g. Fig. 5) also do not indicate the existence of an azeotrope. Brown and Smith (14), who studied the vapor + liquid equilibria of this system at 45", claim the existence of an azeotrope at a mole fraction acetone of and a pressure of mm Hg. In our work, the vapor was found always to be richer in acetone, that is, we did not find an azeotrope. If an azeotrope exists at 45", it must disappear at higher temperatures. Thermodynamic analysis of this system was carried out in two parts. Data reduction for the isotherms at 150 and 200" were carried out using a one-parameter model for the excess Gibbs energy with q,(l) = 0. The results for 150" are shown in Table 4. The Table shows that yi -> 1 as xi -> 1 for this system. Figure 1 1 is a graphical representation of the activity coefficients as a function of mole fraction of acetone in the liquid phase. The differences between the experimental values for the activity coefficients and their calculated values are small; in other words, this

8 CANADIAN JOURNAL OF CHEMISTRY. VOL. 48, 1970 TABLE 3 Thermodynamic analysis of the system CHCI, (1) -acetone (2) at 160 "C* -- Volume fraction of component - - Poynting 2 in the P VI VZL correction y2 liquid phase (atm) (cc/mole) (cc/mole) for yl Poynting. correction for y2 ylpo(exp) Ave. deviation Ave. deviation *Correction to geometric mean K12= 0.010; ~T~~/(Tc, reference fugacity 2 = atm; a12, interaction constant of 1 and 2 = mole/cc W Z U ; om- 0 > Z TC2) = 0.008; 2vI2/(VCI + VC2) = 0.035; reference fugacity 1 = atm MOLE FRACTION ACETONE IN LIQUID RULSE Activity coefficients us. mole fraction acetone in liquid phase of the systems chloroform-acetone at 160 OC. system responds better to the solution model of Chueh and Prausnitz (16). This system is nearly ideal at 150, since the activity coefficients are very nearly equal to one, but at 200" the deviations from ideality are larger. The deviations are positive at 150" and negative at 200". Data reductions for the system at 250 and 270" were carried out using a two parameter dilated van Laar model for the excess Gibbs energy and the unsymmetric convention of normalization for activity coefficients. As can be seen from Table 5, the program calculates f,(')/x, us. x, for each point and, by extrapolating the plot of In f2(')/x2 US. x2 (Fig. 12), Henry's constant of (2) in (1) is obtained, Hz(,). The program also prints out the saturation pressure of component 1, liquid partial molal volumes of both components at infinite dilution, the reference fugacity of component 1 at zero pressure, the molal volumes of the saturated liquid mixture, for each concentration, and the corrected reduced temperatures of the liquid mixture for each concentration (Table 5). This information is then used to evaluate the self-interaction constant, a,,(,, and the dilation constant q2(l) of the dilated van Laar model. Benzene (I) - Carbon Tetrachloride (2) The plot of Tm us. composition (Fig. 3) for this system deviates only slightly from straight line

9 - CAMPBELL AND MUSBALLY: VAPOR PRESSURE AND VAPOR-LIQUID EQUILIBRIA TABLE 4 Thermodynamic analysis of the system CCl, (1) -acetone (2) at 150 OC* -~ ~~~-~.-..--~A~---~ A- Volume fraction of component in the P VI VzL Poynting correction Poynting correction X 2 y2 liquid phase (atm) (cclmole) (cclmole) for yl for y2 ylpo(exp) yzp0(exp) OO OO Ave. deviation Ave. deviation 'Correction to geometric mean K,, = 0.01 ; ~ T,~/(T~ + TCZ) = ; 2vL2/VC1 + VCZ) = ; reference fugacity 1 = 6.07 atrn; reference fugacity 2 = atm; all, interaction constan:of 1 and 2 = mole/cc. MOLE FFACTION ACETONE IN LIOUIO FXASE FIG. 11. Activity coefficients us. mole fraction acetone in liquid phase for the system carbon tetrachlorideacetone at 150 "C. MOLE FRACTION ACETONE FIG. 12. Plot of Henry's constant H2(~+". behavior. The vapor F! liquid equilibrium composition curves (Fig. 6) do not suggest the existence of an azeotrope, since the vapor is always richer in carbon tetrachloride than the liquid. Since the volatilities of the two pure components are very similar, no dilation constant is required and the binary system can be treated, usinga one-parameter model for the excess Gibbs energy. Therefore, data reduction for this system was carried out using the symmetric convention for activity coefficients. The thermodynamic analyses for 150, 200, 250, and 270" are to be found in Depository but that for 200" is given as Table 6.

10 TABLE 5 Thermodynamic analysis of the system carbon tetrachloride (1) - acetone (2) at 250 "C* Molal volume 2 Saturation of saturated - - Corrected pressure of P liquid mixture VI VzL TR of liquid 2 (at,) I (atm) X2 xz Y2 (at m) (cc/mole) $1 $2 (cc/mole) (cc/mole) mixture - 8 c v Vol. fraction cl of component Poynting Poynting 8 2 in the correction correction z xz liquid phase for YI for yz yl "(exp) yzpo(exp) yl p"(calcd) yl (exp) - y,(calcd) yz P"(calcd) yz(exp) - y,(calcd) F Ave. deviation Ave. deviation P m 'Henry's constant at the saturation pressure of solvent determined graphically from above data in upper table = atm. Reference fugacity 1 = atm; saturation pressure 1 = atm; liquid partial molal volume 1 at infinite dilution = cclmole; Vzw = cclmole; azz(,,. self-interaction constant of molecules 2 in the environment of molecules 1 = molelcc; nz(,,, dilation constant of solute 2 in solvent 1 = ; Henry's constant at zero pressure = atm. o n P 2

11 CAMPBELL AND MUSBALLY: VAPOR PRESSURE AND VAPOR-LIQUID EQUILIBRIA TABLE 6 Thermodynamic analysis of the system C6H6 (1) - CCI, (2) at 200 "C* Volume fraction of component - - Poynting Poynting 2 in the P VI VZL correction correction xz y2 liquid phase (atm) (cc/mole) (cc/mole) for y, for y2 ylpo(exp) yzp0(exp) Q Ave. deviation Ave. deviation *Correction to geometric mean K,, = 0.010; 2r1~/(T,~ reference fugacity 2 = atm; a,,, interaction constant of 1 and 2 = rnolelcc. + Tc2) = ; 2v12/(Vc, + VC2) = ; reference fugacity 1 = 11.62atm; MOLE FRACTION OF CARBON IElRACHLORlDE IN LIOUIO PHASE FIG. 13. Activity coefficients us. mole fraction carbon tetrachloride in liquid phase for the system benzene - carbon tetrachloride at 200 OC. The deviation of the activity coefficients from unity is not large but the system cannot be described as ideal. Since this system consists of two non-polar components, it would seem to be the ideal one whereby to test the solution model of Chueh and Prausnitz (16). Table 6 shows that the differences between experimental and calculated values of the activity coefficients are very small. An example of the plot of activity coefficients us. mole fraction carbon tetrachloride (for 200") is given as Fig. 13. Coinparison of Systems The system acetone-chloroform shows negative deviation from Raoult's law and Campbell and Kartzmark (4) have shown that a 1 : 1 compound exists in the solid (and to some extent in the liquid) state. Many theoretical and semitheoretical treatments of this system have been attempted (23, 24). Although there is no doubt of the hydrogen bonding, at room temperature, as indicated by the results of ultraviolet spectroscopy (24) and proton magnetic resonance (25), the attempt to construct a theory is complicated by the fact that both acetone and chloroform are associated in the pure state. Our results (e.g. Table 3) show that the one parameter model of the van Laar equation is not successful in reproducing activity coefficients of the system acetone-chloroform at temper-

12 3184 CANADIAN JOURNAL OF CHEMISTRY. VOL. 48, 1970 atures 150, 160, 170, and 180"; that is, at temperatures at which both components can exist in the liquid state. The dilated van Laar model of Chueh and Prausnitz (16) is based on the assumption that the non-ideality of liquid mixtures is due to the interaction of solute molecules with each other in the environment of the solvent molecules and not to interaction between solute and solvent molecules. It is doubtful if such a concept applies to the acetone-chloroform system. Furthermore, Prausnitz and Chueh's theory applies mostly to non-polar or slightly polar components, but both acetone and chloroform are polar compounds. Neither does the solution model of Prausnitz and Chueh apply in strictness to the system acetone - carbon tetrachloride. Table 4 shows that the one-parameter model of the van Laar equation is fairly successful in reproducing activity coefficients for the temperatures 150 and 200, but the two-parameter model applied at temperatures where one of the components (acetone) is supercritical (Table 5) is apparently not very successful. The same conclusion was reached by Campbell and Chatterjee (1) who used the two-parameter model for the system acetonebenzene at temperatures higher than 200". The system benzene - carbon tetrachloride consisting of two non-polar components should be the ideal system for testing the solution model of Prausnitz. It has been shown, however (26), that the deviation of physical properties from the mixture rule in this system can be ascribed to the association of carbon tetrachloride. Our results (cf. Table 6) show that the oneparameter solution model gives thermodynamically consistent results for the activity coefficients at 150, 200, 250, and 270, but we are at a loss to explain why the activity coefficients are so far removed from unity. Interaction between benzene molecules and carbon tetrachloride molecules might constitute an explanation if there were any reason to suppose such an interaction. It is interesting that, according to Ho, Boshko, and Lu (27), at lower temperatures this system shows very slight positive deviations from Raoult's law but there does seem to be a tendency for these to become less positive as the temperature rises. 1. A. N. CAMPBELL and R. M. CHATTERJEE. Can. J. Chem. 47, 3893 (1969); 48,277 (1970). 2. J. S. ROWLINSON. Liquids and liquid mixtures. Butterworths Scientific Publ., London A. N. CAMPBELL, E. M. KARTZMARK, and H. FRIESEN. Can. J. Chern. 39, 735 (1961). 4. A. N. CAMPBELL and E. M. KARTZMARK. Can. J. Chem. 38, 652 (1960). 5. J. V. ZAWIDSKI. Z. phy. Chern. 35, 129 (1900). 6. H. ROCK and W. SCHR~DER. Z. phy. Chern. (Frankfurt) 11, 41 (1957). 7. A. N. CAMPBELL, E. M. KARTZMARK, and R. M. CHATTERJEE. Can. J. Chem. 44, 1183 (1966). 8. S. YOUNG. Distillation principles and processes. MacMillan, London G. SCATCHARD, S. E. WOOD, and J. M. MOCHEL. J, Amer. Chem. Soc. 70, 1723 (1948). 10. R. T. FOWLER and S. C. LIM. J. Appl. Chem. 6, 75 (1956). 11. A. 'N. CAMPBELL and W. J. DULMAGE. J. Arner. Chern. Soc. 70, 1723 (1948). 12. G. C. GERRITS. Proc. Acad. Sci. Amsterdam. 7, 167 (1904). 13. W. H. SEVERNS, A. SESONSKE, R. H. PERRY, and R. L. PIGFORD. J. Arner. Inst. Chern. Eng. 1, 401 (1955). \ - ~ - -,. 14. I. BROWN and F. SMITH. Aust. J. Chem. 10, 423 (19571 \---.,. 15. S. R. M. ELLIS and L. M. ROSE. Birmingham Univ. Chem. Eng. 14, (2) 45 (1963). 16. P. L. CHUEH and J. M. PRAUSNITZ. Ind. Eng. Chern. 60, 35 (1968); Computer calculations for high pressure vapor liquid equilibria. Prentice-Hall H. C. VAN NESS. Classical thermodynan~ics of non-electrolyte solutions. Pergamon Press, New York P. L. CHUEH, N. K. MUIRBROOK, and J. M. PRAUS- NITZ. Arner. Inst. Chem. Eng. J. 11, 1097 (1965) REDLICH and A. T. KISTER. Ind. Eng. Chern. 40, 345 (1948). 20. E. F. G. HERINGTON. Nature, 160, 610 (1947). 21. W. SWIETOSLAWSKI and A. KREGLEWSKI. Bull. de 1'Acad. Polonaise des Sci. Sciences C1 111-Vol. 11, No. 4 (1954), J. P. KUENEN and W. G. ROBSON. Phil. Mag. 3, 149, 622 (1902). 23. J. A. BARKER and F. J. SMITH. J. Chern. Phys. 22, 375 (1954). 24. A. MUNSTER. Trans. Faraday Soc. 46, 165 (1950). 25. C. M. HUGGINS, G. C. PIMENTEL, and J. N. SHOOLERY. J. Chem. Phys. 23, 1244 (1955). 26. A. SCHULZE. Z. phy. Chern. 86, 309 (1913). 27. J. C. K. Ho, 0. BOSHKO, and B. C. Y. Lu. Can. J. Chem. Eng. 39, 205 (1961).