V. FRIED,D. R. FRANCESCHETTI, AND A. S. GALLANTER
1476
The equilibrium constant Kn2for the reaction AgCl(s)
+ C 1 - 5 AgC12-
is also simple to interpret in terms of solvation. The more strongly solvated is the complex ion AgC1,- and the less strongly solvated is C1-, the larger is this constant. For DMSO this constant falls in the same range as it does for propylene carbonate2 (log K,z = 1.00) dimethylf~rmamide~ (log KeZ= 1.80),nitroethanes (log K,z = l . l ) , acetone6 (log K,, = 0.3), and acetonitrile6 (log K,2 = 0.2). On addition of water (Figure 4), KBzdecreases by more than three orders of magnitude and appears to achieve a plateau a t ca. 10-2.0. For pure water it
reaches the va1ueI6 log Kat = -4.7. The behavior of K,z in water containing small abounts of DMSO has not been investigated, but it may be as dramatic as that in DMSO containing small amounts of water. This pronounced shift results both from the decreased solvation of silver by DMSO (as reflected also by K,1) and from the increased solvation of chloride by hydrogen bonding to the water. Acknowledgment. This work was supported by the Air Force Cambridge Research Laboratories, Office of Aerospace Research, under Contract No. 19 (628) -6131, but does not necessarily constitute the opinions of that agency. The authors thank Dr. N. A. Rumbaut for graciously providing them with unpublished data.
Thermodynamic Properties of the System Carbon Tetrachloride-Tetrachloroethylene by V. Fried, D. R. Franceschetti, and A. S. Gallanter Department of Chemistry, Brooklyn College of the City University of New York, Brooklyn, New York (Received October 14, 1 9 6 8 )
11210
We have proved experimentally that the system carbon tetrachloride-tetrachloroethylene is nearly ideal. The equilibrium pressure as a function of the mole fraction of carbon tetrachloride, 2,is given by P 144.4 476.62 at 70’ and P = 95.1 348.32 at 60’. Similarly, the molar volume is expressed by V = 102.731 - 5.6112. This indicates that the interaction energy between unlike molecules is the arithmetic mean of the interaction energies for like molecules.
+
Introduction Extensive investigations of charge-transfer interactions in various pure solvents have been reported in recent years.’J Little study has been made of molecular complex formation in solvent mixtures. Our present research, the study of some electron donor-acceptor interactions in solutions requires knowledge of the thermodynamic behavior of our solvent mixture, carbon tetrachloride-tetrachloroethylene. Carbon tetrachloride forms molecular complexes with some electron donors. Its coordination products with benzene and benzene derivatives are well known.*-’ The R electrons of the benzene ring interact with the electron-deficient chlorine atoms in the carbon tetrachloride molecule. The R electrons in tetrachloroethylene are less available for complex formation because of the electron-withdrawing effect of the halogens directly attached to the olefinic carbons. It is well known that while ethylene acts as an electron donor, and tetracyanoethylene as a strong electron acceptor, tetrachloroethylene is practically nonreactive.’I2 The Journal of Physical Chemistry
+
+
To find out if there is any kind of interaction between the two different molecules of the solvent mixture, carbon tetrachloride-tetrachloroethylene, we have measured the thermodynamic properties of mixing in this system. Experimental Section Chemic& Carbon tetrachloride and tetrachloroethylene were purified as described earlier.*v9 The (1) G. Briegleb, “Molekulverbindungen und Koordinationsverbindungen in Einsel darstelungen. Electronen-DonatorAcceptor Komplexe,” Springer-Verlag, Berlin, 1961. (2) L. J. Andrew6 and R. M. Keefer. “Molecular Complexes in Organic Chemistry,” Holden-Day, San Francisco, Calif.. 1964. (3) R. P. Rastogi, J. Nath, and J. Misra, J . Phys. Chem., 71, 1277 (1967). (4) R. P. Rastogi and R. K. Nigam, Trans. Faraday SOC.,55. 2005 (1959). (5) J. B. Ott, J. R. Goates, and A. H. Budge, J . Phys. Chem., 66, 1387 (1962). (6) J. R. Goates, R . J. Sullivan, and J. B. Ott. tbid., 63, 589 (1959). (7) R. Anderson and J. M. Prausnits, J. Chem. Phys., 39, 1226 (1963). (8) V. Fried, D,R. Franceschetti, and G. B. Schneier, J. Chem. Eng. Data, 13, 416 (1968). (9) V. Fried, P. Gallant, and G. B. Schneier, (bid., 12, 504 (1967).
THERMODYNAMIC PROPERTIES OF CARBON TETRACHLORIDE-TETRACHLOROETHYLENE measured densities of both components at 25' (1.58413, 1.61442 g/cc) and the normal boiling points (76.73', 121.03') were in good agreement with the literature values.lo At atmospheric pressure the boiling points and condensation temperatures of both substances, measured with a Beckmann thermometer in a modified Swietoslawski differential ebulliometer," differed by only 0.002 and 0.004°, respectively. Both these results and the cooling curve measurements indicate that both samples did not contain impurities having any significant effect on the measured thermodynamic properties of mixing. Tetrachloroethylene was stabilized with thymol. Apparatus and Procedure. The vapor-liquid equilibrium data were obtained with a modified Gillespie circulation equilibrium stiLil A Swietoslawski-type ebulliometer" filled with deionized and distilled water was connected in parallel with the system; from the boiling point of water, the corresponding pressure of the system was determined (within 0.1 mm). This arrangement eliminates even the errors in the boiling point measurements owing to changes in the barometric pressure during the experiment. The temperature of the equilibrium phases and the boiling point of water were measured with mercury thermometers calibrated by the National Bureau of Standards. The inaccuracy in the temperature reading was less than 0.02'. We tried several methods to analyze the equilibrium vapor and liquid samples. Because of the similarity in physical properties of CCL and C2C14, none of the methods gave perfect results. The refractive index measured with a Bausch and Lomb refractometer gave the best results. The values found for known compositions at 25 f 0.02' are given in Table I. The calibration points were fitted by the method of least squares to the equation nZbD=
a
+ bx + cx2
(1)
Table I: Refractive Index As a Function of the Composition: System Carbon Tetrachloride-Tetrachloroethylene at 25" ZCCl4
O.OO0 0.094 0.121 0.163 0.214 0.286 0.322 0.345 0.480 0.485 0.506 0.548 0.559 0.649 0.667 0.668
~ W I
1.5034 1.4994 1.4983 1.4965 1,4943 1.4910 1.4896 1.4885 1.4824 1.4822 1.4813 1.4794 1.4787 1.4745 1.4738 1.4737
ZCCI4
n26D
0.700 0.729 0.748 0.781 0.771 0.790 0.805 0.824 0.843 0.849 0.880 0.883 0.915 0.953 0.958 1.OOO
1.4722 1.4708 1.4699 1.4695 1.4688 1.4680 1.4672 1.4663 1.4654 1.4650 1.4635 1.4634 1.4618 1.4600 1* 4597 1.4577
1477
where x is the mole fraction of carbon tetrachloride. The constants a, b, and c have the following values: 1.50340, -0.04173, and 0.00405. The standard deviation, =t9 X lo", is within the limit of the experimental error. An inaccuracy in the refractive index measurements of fO.OOO1 results in an inaccuracy of f0.0025 in the mole fraction. Considering the error introduced by evaporation of the sample during the measurement, the over-all absolute error in the composition was no more than f0.4 mol %. This inaccuracy affected the equilibrium pressure by about =t2 mm at 70' and about f 1 . 3 mm at 60'. Single-stem pycnometers of about 10-ml capacity were used for density measurements.
Results and Discussion The experimental vapor-liquid equilibrium data a t 60 and 70' are presented in Table 11. In calculating the activity coefficients, yi, the authors assumed that the vapor phase is an ideal solution of nonideal vapors. The coefficient AE, defined by the equation Ai2
= 2B12
- Bii - B22
(2)
is zero and the following equation was used
(3) B22, BIZare the second virial coefficients, yi and X i the mole fractions in the vapor and liquid phase, respectively, Pio the vapor pressure of the pure component, and P the total equilibrium pressure. The correction term, Cil is given by the equation
&I,
Ci
=
exp[(P
- Pi") (Bii - Vil")/RT]
(4)
where ViP is the molar volume of pure component i in the liquid phase and R is the ideal gas constant. As is evident from Table 11, Ci appreciably affects the value of the activity coefficient. As the mole fractions in the liquid phase and vapor phase are both affected by the same experimental error, k 0 . 4 mol %, the uncertainty in the determination of the activity coefficients is about 2%. All the data necessary to calculate the activity coefficients are collected in Table 111. The virial coefficients of the pure components were calculated from Pitzer's equation12 using acentric factors of 0.230 and 0.189 for tetrachloroethylene and carbon tetrachloride, respectively. The uncertainty in Ci due to allowed 10% inaccuracy of Bii is insignificant. The activity coefficients indicate that the system CClrC2C14 behaves ideally within our experimental (IO) J. Timmermans, "Physico-Chemical Constants of Pure Organic Compounds." Elsevier, New York, N. Y.,1950. (11) E. H&la,J.'Pick, V. Fried, and 0. Villm, "Vapor-Liquid Equilibrium," Pergamon Press, London, 1967. (12) K. 8. Pitzer and R. F. Curl, J. Amer. Chem. Soc., 79, 2369 (1967). Volume Y3,Number 6 Mau 1969
V. FRIED,D. R, FRANCESCHETTI, AND A. S. GALLANTER
1478
Table 11: Vapor-Liquid Equilibrium: System Carbon Tetrachloride-Tetrachloroethylene ax14
WCI4
P
CCCIl
60
0.049 0.068 0.152 0.175 0.231 0.242 0.296 0.317 0.391 0.416 0.451 0.480 0.523 0.564 0.577 0.657 0.737 0.778 0.871
0.188 0.253 0.452 0.490 0.578 0.588 0.660 0.680 0.749 0.765 0.789 0.812 0.830 0.858 0.861 0.900 0.927 0.941 0.970
112.8 118.9 147.9 155.9 175.3 180.0 198.5 205.4 231.6 238.7 252.6 264.2 276.6 292.2 297.4 321.6 352.8 364.0 399.3
1I020 1.020 1.018 1.017 1.016 1.016 1.015 1.014 1.013 1.012 1.011 1.011 1.010 1.009 1.008 1.007 1.005 1.005 1.003
70
0.043 0.099 0.180 0.220 0.275 0.361 0.425 0.488 0.582 0 620 0.674 0.711 0.878 0.935
0.160 0.317 0.483 0.545 0.618 0.708 0.763 0.808 0.852 0 873 0.896 0.913 0.968 0.984
165.6 192.6 231.1 249.5 275.0 314.4 346.9 376.6 423.6 442.7 465.9 485.0 563.4 590.8
1.025 1.023 1.021 1.020 1.019 1.017 1.015 1.013 1.011 1.010 1.008 1.007 1.003 1.001
Temp, OC
I
I
CCCzlk
YCC Ik
YCzoIl
0.998 0.998 0.995 0.995 0.993 0.992 0.991 0.990 0.988 0.987 0.986 0.985 0.984 0.983 0.982 0.980 0.977 0.976 0.973
1.00 1-01 1.01 1.00 1.00 1-00 1.01 1.01 1.01 1.oo 1.01 1.02 1.oo 1.00 1.oo 1.oo 1.00 1.00 1.00
1.01 1.00 1.oo 1.01 1.01 1.02 1.00 1.00 0.99 1.oo 1.01 0.99 1.02 0.99 1.03 0.97 1.01 1.00 0.98
0 998 0.996 0.993 0.991 0.989 0.986 0.984 0 981 0.978 0.976 0.975 0.973 0.967 0.964
1.02 1.02 1.03 1.02 1.02 1.02 1.02 1.02 1.02 1.02 1.01 1.02 1.01 1.01
1.03 1.03 1.02 1.02 1.01 1.00 1.00 0.98 1.03 1.02 1.02 1.00 1.01 1.00
I
I
Table 111: Physical Properties of the Pure Substances" Temp,
O O
60 70 a
V2,
Vi, ml mol-1
Pio, mm
-Bii, ml mol-1
ml mol-1
Pao, mm
101.456 102.812
444.16 617.07
1139 1059
106.571 107.679
94.88 141.38
-B1& ml mol-1
1737 1592
Subscript 1, CCI,; subscript 2, C2CId.
error. Equations
+ bz = 144.4 + 4 7 6 . 6 ~(at 70') P = a + bz = 95.1 + 348.33 (at 60')
P and
=a
(5)
(6) affirm the near ideality of the system. The deviation of the constants in eq 5 and 6 from those calculated from the vapor pressures of the pure components, a = P2' = 141.4 mm and b = (Plo- Pa0)= 475.7 mm at 70' and a = P2' = 94.9 mm and b = (PI' - P2') = 349.3 mm a t 60', and also the standard deviations h1.77 and f 1 . 0 3 mm are within the limits of experimental error. The Journal of Phyaical Chemietry
Table IV presents the densities and molar volumes as a function of the composition at 25'. The linear equation
V
=a
+ bz = 102.731 - 5.6112
(7)
again proves the system to be nearly ideal. The deviation of the constants in eq 7 from those calculated from the volumes of the pure components, a = V2' = 102.718 ml and b = (Vlo - V2') = 5.614 ml, aswellas the standard deviation f0.024 ml mol-', is within the limits of experimental error. The relatively low values of the heat of mixing found
THERMODYNAMIC PROPERTIES OF CARBON TETRACHLORIDE-TETRACHLOROETHYLENE Table IV: Densities and Molar Volumes As Function of the Composition: System Carbon Tetrachloride-Tetrachloroethylene at 25"
1479
and log yz
=
0.01115
xz-to
Similar values were found a t 60'. With the standard state as the pure component a t the temperature and pressure of the system, these limiting values are also the 102.718 0.0000 1.61.442 maximum values of the activity coefficients. This 1.61036 101.964 0.1359 100.880 0.3321 1.60430 semiempirical theory thus predicts the nearly ideal 0 4825 1.69976 100.037 behavior of the CC14-C2CLsystem very well. 1,59283 98.752 0.7110 The system CCl4-C2CI4is found experimentally to be 97.104 1.0000 1.58413 nearly ideal. The interactions between the two unlike molecules are therefore negligibly small. Basing our judgment only on the size of the two different molecules, we expect this behavior (VP/V," = 0.952 a t 60' and by Poon and Lu13 also affirm the near ideality of the VI"/VZ' = 0.955 at 70'). Nevertheless, considering system. the symmetry and shape of both molecules, the ideal Of the theories of s o l ~ t i o n s ' ~tested - ~ ~ only the regular behavior does not seem too obvious. The CCh molecule solution theory predicts the near ideality of the system. is spherically symmetrical since the sp*-hybridized The values given by some of these theories are so much carbon is bonded to four identical substituents while the in error that it is not worthwhile to mention them here. CzCl4 molecule, containing trigonally hybridized As can be seen from the difference in the solubility carbons, is planar. The C-C1 bond distance is shorter parameters (6' - 62 = -0.74 at 60°), the solubility in the olefin halide than in carbon tetrachloride. I n parameter theory18 predicts a small deviation from spite of it the system is nearly ideal. This can be ideal behavior. The value of the excess Gibbs free in that the reactivity of the C-C bond in rationalized energy, calculated from this theory for z = 0.5, is tetrachloroethylene is already reduced by the with59 J mol-', which is in very good agreement with the drawal of electrons and that it is also buried by four measured heat of mixing13 and also with our data, large chlorine atoms and therefore unavailable for intertaking into consideration the 2% uncertainty. actions. It is also significant that the electrons in the Very good results were also obtained using the semiempirical treatment of solutions proposed by E r d o ~ . ' ~ ?r bond are located in thick layers above and below the u bond, and this brings the molecule closer to spherical Based on this theory, the following equation may be symmetry. written for the limiting values of the activity coefficients d
WCl4
V
I
lim log y1 = z1-0
lim log yz = xa-rO
1 2.303R T
1 2.303R T
where Ucl = 6580 J mol-' and Ucz = 8270 J mol-' are the cohesion energies a t 7OoZ0 and [ P J = 224.2 and [Pz] = 247.3 are the parachors a t the same temperature.21 The limiting values of the activity coefficients are log 71 = 0.01030 x1-0
Acknowledgment. This investigation was supported in part by Undergraduate Science Education Program Grants GY-2860, National Science Foundation. The authors are indebted to Dr. I. A. Kaye for many useful discussions during the work. (13) D. P. L. Poon and B. C. Y. Lu, J . Chem. Eng. Data, 13, 435 (1968). (14) P. J. Flory, J . Amer. Chem. Soc., 8 7 , 1733 (1965). (15) I. Prigogine, A. Bellemans, and A. E. Charles, J . Chem. Phys., 24, 518 (1965). (16) I. Prigogine, "Molecular Theory of Solutions," North-Holland Publishing Do., Amsterdam, 1957. (17) E. A. Guggenheim. "Mixtures." Clarendon Press, Oxford, 1952. (18) J. H. Hildebrand and R . L. Scott, "The Solubility of Nonelectrolytes," Dover, New York. N. Y., 1964. (19) E . Erdas, Collect. Czech. Chem. Commun., 21, 1528 (1956). (20) V. Fried and G. B. Schneier, J . Phys. Chem., 7 2 , 4688 (196s). (21) R. R. Dreisbach, "Physical Properties of Chemical Oompounds11," American Chemical Society, Washington, D. C., 1965.
Volume YS, Number 6 M a y 1960
