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Thermodynamic excess functions of methanol + piperidine at 298.15 K
J. Chem. Thermodynamics1975,
I,
1125-l 130
Thermodynamic excess functions of methanol + piperidine at 298.15 K KOICHIRO HIDEKAZU
NAKANISHI, TOUHARA
HIDE0
WADA,
and
Departmentof Industrial Chemistry,Kyoto University, Kyoto 606, Japan (Received2I April 1975) Vapour pressures of methanol + piperidine at 298.15, 308.15, and 318.15 K were measured by a static method. Excess enthalpies and densities of the same mixtures at 298.15 K were also determined with an isothermal dilution calorimeter and a pyknometer. The excess functions were evaluated from these results; the values for mole fraction x = 0.5 at 298.15 K are: GE = -834.1 J mol-I, HE = -3159.1 J mol-I, TSE = -2325.0 J mol-I, VE = -l.26cm3 mol-I.
1. Introduction As a part of a thermodynamic study of alcohol + amine mixtures,(r.‘) the excess Gibbs free energies GE, the excess enthalpies HE, and the excess volumes vE of methanol (MeOH) + piperidine (C,H,,N) at 298.15 K have been determined. The purpose of the present report is to evaluate the influences of ring structures and aromaticity on the thermodynamic properties of binary mixtures by comparing the results in the present mixtures with those for methanol + non-cyclic aliphatic amine and methanol f pyridine.
2. Experimental MATERIALS
Methanol, a spectrograde reagent, was used without further purification. Piperidine, a guaranteed reagent, was distilled from CaCl, under dry nitrogen. As shown in table 1, the densities and vapour pressures of these liquids agreed well with literature values.(3-s) VAPOUR-PRESSURE MEASUREMENT The vapour pressures of mixtures were determined by a static method. The design of our vapour-pressure apparatus(g) was based upon one described by Hermsen and Prausnitz.(“) However, several modifications were made in the degassing device, the sample vessel, and the manometer. The sample vessel was made of Pyrex glass and its
1126 TABLE
K. NAKANISHI,
H. WADA,
AND H. TOUHARA
1. Densities p, refractive indices nD, vapour pressures p, and boiling temperatures pure fluids (Torr = (101.325/760) kPa) This work
Tb of
Literature
Methanol
p(298.15 K)/g crnw3 p(298.15 K)/Torr
0.78673 126.90
0.78664 cm 126.89 c4) 127.21 G) 126.96 (W
Piperidine
p(298.15 K)/g crnT3 nD(298.15 K) p(298.15 K)/Torr Tb(760 To&/K
0.85691 1.4505 30.49 379.45
0.85674 U) 1.4501 (3) 30.19 @) 379.55 (a
volume was about 50 cm3. The null manometer was a U-shaped glass tube. The inner diameter of the part of this tube used to read mercury levels was 14 to 15 mm. Both components were mixed in a glass vessel and degassed by repeating freezeand-melt procedure under vacuum. Degassed liquids were then transferred directly into the sample vessel without exposing them to air. Glass needle valves having Teflon needles were used wherever the sample liquids were allowed to pass through them. A Teflon-coated magnet was inserted in the sample vessel. It was driven by a submarine magnetic stirrer placed underneath the vessel and was proved effective in shortening the time required to attain vapour + liquid equilibrium. During the pressure measurements, the temperatures of the sample vessel and the null manometer immersed in a double thermostat were controlled within kO.01 K. The position of the thermostat was adjusted by lifting it with a laboratory jack. The total pressures of sample liquids at three different temperatures were read on mercury manometers by means of a cathetometer and a travelling microscope. Capillary corrections were applied to the manometer readings and the vapour pressures were then calculated by a conventional method using the local value of the acceleration of free fall and the density of mercury. The temperature of the thermostat was monitored by a Beckmann thermometer which was recalibrated against a calibrated standard thermometer. Its precision was +0.03 K. The vapour-pressure measurement was first done at 298.15 K, then the temperature of the thermostat was raised successively to 308.15 and 318.15 K, and the vapour pressure was measured at each temperature. The temperature was lowered again to 298.15 K and a duplicate measurement was made. Agreement with the first measurement was confirmed to complete one series of measurements. Equilibrium compositions were determined after the pressure measurements either by density or refractive index measurements. The densities of the sample mixtures were determined with a bicapillary pyknometer, the volume of which was about 30 cm3, and the refractive index measurements were made with a Pulfrich refractometer. Except for the piperidine-rich region where the density of the mixture goes through a flat maximum, the density measurement was preferred on account of
EXCESS
FUNCTIONS
OF
CHeOH
1127
+ CsHllN
better accuracy. The following smoothed values were used to convert the density p or the refractive index n, of a sample mixture at 298.15 K to the composition. P = bwl
+-M2(1 -xl))lwl/Pl)xl +w,IP,)(l --Xl> +x,(1-x,){ -5.O36-1.O25(2x,-l)+O.2771(2x1-1)2}
nn = nlxl+n~(~-x,)+x,(l-x,){0.1290+0.0578(2x,-1)+0.0174(2x,-1)2}. Subscripts 1 and 2 refer to methanol mass and x the mole fraction. EXCESS-ENTHALPY
and piperidine,
g cm-‘],
(1) (2)
respectively; M is the molar
MEASUREMENTS
An isothermal displacement calorimeter was constructed and used to measure the excess enthalpies at 298.15 K. Calibration of the calorimeter for exothermic mixing was done with dichloromethane -I- p-dioxane. The results(l’) agreed quite well with the values reported by Murukami and Benson.(“) To use this calorimeter in the present study, care had to be taken with the following problems. First, the O-ring packing of the Dewar vessel of the calorimeter must be completely resistent to highly corrosive amine-containing mixtures. A Viton E-60C O-ring was found to be suitable and used throughout the measurements. Secondly, the temperature rises on mixing are rather large for these liquids. To avoid the accumulation of errors on successive additions of the second component, the determination of composition must be done by calculations on a mass basis. Conversion of the volume added to the mass of the second component will lead to a non-negligible error inlthe composition. This situation has forced us to change the final design of the second-component feeding device of our calorimeter.(l’)
3. Results VAPOUR
In at to at
PRESSURES
table 2 are recorded the isothermal total pressures for x,MeOH + (1 -x&H, ,N 298.15, 308.15, and 3 18.15 K. Negative deviations from the Raoult’s law are found be fairly large. The errors in the vapour pressures are estimated to be f0.3 Torr 298.15 K and may be slightly larger at 308.15 and 318.15 K.?
EXCESS FUNCTIONS
From the above results, the excess Gibbs free energies GE were calculated by the Barker method.(13) Table 3 shows the values used for the molar volumeV$ and the second virial coefficient Bii of the pure liquids. Slightly different values of B,, are available for methanol(‘* ’ 6, and piperidine. u’) It was found that the calculated values of GE did not change significantly with the value assigned to Bii. There are no data on the cross virial coefficient Bij for this mixture. Therefore, 6, z = 2Bl 2 - B, 1Bz2 was assumed to be zero. t Throughout
this work
Torr
= (101.325/760)
kPa.
1128
K. NAKANISHI,
TABLE
2. Vapour
pressuresp,
H.
WADA,
excess enthalpies Crorr
AND
H.
TOUHARA
HE, and densities
= (101.325/760)
p for xlMeOH
+ (1 - xl)CSHllN
kPa)
p(298.15 ~(308.15 ~(318.15
Xl a K)/Torr K)/Torr K)/Torr
0 30.49 51.05 81.84
0.1196 32.83 56.00 92.1,
0.2058 35.7, 61.27 101.83
0.3002 39.6, 68.26 112.40
0.3868 42.90 73.7, 125.4,
0.4924 51.95 89.0, 149.0,
p(298.15 ~(308.15 ~(318.15
Xl a K).Torr K)/Torr K)/Torr
0.6754 75.43 127.54 208.2,
0.7829 94.1s 156.7, 252.59
0.8422 104.6* 173.6X 278.86
0.8815 110.84 183.56 293.9,
0.9393 120.15 197.6, 315.40
1 126.90 208.8, 333.2,
s”(298.15
K)/J
mol-l
-272.3 0.0309
-607.2 0.0670
-1130.5 0.1359
-1614.9 0.1989
-2032.1 0.2583
-2365.5 0.3091
%(298.15
K)/J
mol-l
-2687.2 0.3670
-2906.2 0.4170
-3047.8 0.4623
-3123.0 0.4940
-3186.3 0.5280
-3204.5 0.5546
G”(298.15
IQ/J
mol-l
-3205.8 0.5779
-3183.9 0.6062
-3107.6 0.6441
-2993.5 0.6775
-3020.0 0.6753
-2702.9 0.7317
%(298.15
K)/J
mole1
-2351.1 0.7791
- 0.8310 1902.4
-1361.6 0.8845
-766.0 0.9372
-310.9 0.9744
X1 p((298.15
K)/g
crnT3
0.1008 0.85773
0.2107 0.85700
0.3026 0.85642
0.3897 0.85502
0.5021 0.85182
Xl ~(298.15
K)/g
cme3
0.6949 0.83976
0.7974 0.82765
0.8571 0.81831
0.9054 0.80961
0.9547 0.79856
0.5992 63.06 107.80 177.41
0.5952 0.84729
4 Strictly speaking these x1 values are valid only at 298.15 K, since the sampling was done that temperature. Corrections to x1 for other temperatures were estimated to be less than 0.0001.
TABLE
3. Molar
volumes
Temperature T/K
Vt and second
virial
coefficients
Bii of the pure
MeOH V?/cm3
298.15 308.15 318.15
The Redlich-Kister
40.7 41.2 41.7
mol - 1
at
liquids
C&LN &/cm”
mol - 1
-1970 -1710 -1490
Vk/cm3
mol - 1
99.4 100.5 101.6
equation was used to express the composition
Bt,/cm3mol-1 -2116 -1910 -1740
dependence of GE :
GE/J mol- ’ = xlxz k A&, -xZ)nW1. n=l The results of iterative calculations indicate that the present results could be fitted satisfactorily by equation (3) with 4 constants, the values of which are given in table 4. The standard deviations o(p) of the vapour pressure are also included in the table. The excess enthalpies at 298.15 K are given in table 2. They were also fitted to an equation of the same type as equation (3) with HE replacing GE; the constants are given in table 4. The standard deviation o(HE) was 9.51 J mol-I. These two functions GE and HE were combined to calculate the excess entropies SE at 298.15 K.
EXCESS JWNCT’IONS OF CH,OH + C~HIIN
1129
TABLE 4. Coefficientsof the Redlich-Kister equation and smoothed values of excessfunctions for x = 0.5 (Torr = (101.325/760)kPa) T/K 298.15 GE
308.15 318.15
HE I/E
298.15
42
Al
-1239.5
-3336.2 -3100.3 -2764.2 - 12636.21
298.15 -5.036 Smoothed valuesfor x = 0.5 and T = 298.15 GE/Jmol - 1 HE/J mol-l -834.1 -3159.1
A
o(plTorr) 0.55 0.69
-1152.1 -1091.0
187.7 305.8 228.8
909.8 805.3 826.5
-3748.52
2038.28
2149.70
-1.025
0.2771
1.11
-
K.
TP/J mol - ’ -2325.0
P/cm3 mol - 1 -1.26
From the densities used to determine the equilibrium compositions, the excess volumes VE/cm3 mole1 were evaluated and fitted to a polynomial equation with three constants given in table 4. The standard deviation o(V’) was 0.014 cm3 mol-‘. Smoothed values of all the excess functions for x = 0.5 are also included in table 4. Figure 1 illustrates the dependence of GE, HE and TSE on x1. All the excess functions are fairly symmetrical with respect to x1, although minima are found at about
x1 = 0.55.
FIGURE 1. Excessfunctions of xlMeOH + (1 - xI)C5H11N at 298.15K. _---_
TSE;
-
XE:
-.-.-*.
HE.
- GE;
4. Discussion The excess functions of methanol
+ piperidine can be summarized as follows:
HE
VE< 0.
This is similar to the pattern observed for methanol + diethylamine or methanol + n-butylamine studied by us previously. (2) Comparing secondary amines, we find that TSE for MeOH + &HI ,N is less negative than that for MeOH + Et,NH, suggesting
1130
K. NAKANISHI,
H. WADA, AND I-I. TOUHARA
significant loss of internal degrees of freedom in E&NH molecule upon mixing. HE is also less negative for MeOH + C5H11N, in spite of the fact that the pK, of piperidine (11.22) is almost the same as that of Et,NH (10.98).(‘*) This implies that the effect of the basicity of the N-atom in amine molecules on the excess functions is of minor importance and that, in MeOH + C+H, ,N, there should be a packing effect which reduces the O-H---N hydrogen bonding energy. This interpretation may be supported by the fact that VE for MeOHf CSHllN is almost half of that for MeOH -I- Et,NH.@) Since the entropy effect mentioned above is predominant, GE for MeOH + C,H,,N is more negative than that for MeOH + E&NH. Several investigations including our unpublished work(lg’ have been made on an alcohol -I- pyridine. The excess functions for MeOH + C,H,N follow the same pattern as that given above. However, the absolute values of GE, HE, and TSE are smaller than those for MeOH f CSHIIN. The presence of n-electron interaction and NH-O hydrogen bonding, which might be responsible for this difference, can be proved to have only minor effects, and it may be concluded that the difference between the excess functions of MeOH -I- C,HSN and those of MeOH + C,H1,N can be ascribed to the weaker basicity (pK, = 5.23) of pyridine.(l*) The cooperation of Mr M. Kitamura in the construction of the vapour pressure apparatus is gratefully acknowledged. We also thank Professor N. Watanabe for stimulating discussions. REFERENCES
1. Nakanishi,K. ; Shirai, H. Bull. Chem. Sot. Japan 1970,43,1634. 2. Nakanishi,K.; Touhara,H.; Watanabe,N. Bull. Chem. Sot. Japan 1970,43,2671. 3. Riddick, J. A.; Bunger,W. B. Organic Solvents. Wiley: New York. 1970. 4. Manufacturing Chemists’Association,Data Sheets,table 23-2-1-(1.1000)-k, Thermodynamic Research Center,TexasA & M Univ., CollegeStation,Texas,1960. 5. Gibbard,H. F.; Creek,J. L. J. Chem. Eng. Data 1974,19, 308. 6. Polak, J.; Lu, B. C.-Y. J. Chem. Eng. Data 1972, 17,456. 7. Timmermans, J. Physico-Chemical Constants of Pure Organic Compounds. Elsevier:New York. 1965. 8. Osborn,A. G.; Douslin,D. R. J. Chem. Eng. Data 1968, 13, 534. 9. Wada,H. M. Eng. Thesis,Kyoto Univ., 1973. 10. Hermsen,R. W.; Prausnitz,J. M. Chem. Eng. Sci. 1963,18,485. 11. Touhara,H.; Ikeda, M.; Nakanishi,K.; Watanabe,N. J. Chem. Thermodynamics 1975,7,887. 12. Murakami, S.; Benson,G. C. J. Chem. l%ermodynamics 1969,1,559. 13. Barker,J. A. Aust. J. Chem. 1953,6,207. 14. Freydank,H.; Ratzsch,M. 2. Phys, Chem. (Leipzig) 1971,248,83. 15. Scott, D. W. J. Chem. Thermodynamics 1971, 3, 649. 16. Kudchadker,A. P.; Eubank,P. T. J. Chem. Eng. Data 1970, 15, 7. 17. Cabani,S.; Ceccanti,N. J. Chem. Thermodynamics 1973,5, 9. 18. Tamres,M.; Searles,S.; Leighly,E. M.; Mohrman,D. W. J. Amer. Chem. Sot. 1954, 76, 3783.
19, Nakanishi,K.; Ashitani,K.; Touhara,H. Unpublishedwork.
I,
1125-l 130
Thermodynamic excess functions of methanol + piperidine at 298.15 K KOICHIRO HIDEKAZU
NAKANISHI, TOUHARA
HIDE0
WADA,
and
Departmentof Industrial Chemistry,Kyoto University, Kyoto 606, Japan (Received2I April 1975) Vapour pressures of methanol + piperidine at 298.15, 308.15, and 318.15 K were measured by a static method. Excess enthalpies and densities of the same mixtures at 298.15 K were also determined with an isothermal dilution calorimeter and a pyknometer. The excess functions were evaluated from these results; the values for mole fraction x = 0.5 at 298.15 K are: GE = -834.1 J mol-I, HE = -3159.1 J mol-I, TSE = -2325.0 J mol-I, VE = -l.26cm3 mol-I.
1. Introduction As a part of a thermodynamic study of alcohol + amine mixtures,(r.‘) the excess Gibbs free energies GE, the excess enthalpies HE, and the excess volumes vE of methanol (MeOH) + piperidine (C,H,,N) at 298.15 K have been determined. The purpose of the present report is to evaluate the influences of ring structures and aromaticity on the thermodynamic properties of binary mixtures by comparing the results in the present mixtures with those for methanol + non-cyclic aliphatic amine and methanol f pyridine.
2. Experimental MATERIALS
Methanol, a spectrograde reagent, was used without further purification. Piperidine, a guaranteed reagent, was distilled from CaCl, under dry nitrogen. As shown in table 1, the densities and vapour pressures of these liquids agreed well with literature values.(3-s) VAPOUR-PRESSURE MEASUREMENT The vapour pressures of mixtures were determined by a static method. The design of our vapour-pressure apparatus(g) was based upon one described by Hermsen and Prausnitz.(“) However, several modifications were made in the degassing device, the sample vessel, and the manometer. The sample vessel was made of Pyrex glass and its
1126 TABLE
K. NAKANISHI,
H. WADA,
AND H. TOUHARA
1. Densities p, refractive indices nD, vapour pressures p, and boiling temperatures pure fluids (Torr = (101.325/760) kPa) This work
Tb of
Literature
Methanol
p(298.15 K)/g crnw3 p(298.15 K)/Torr
0.78673 126.90
0.78664 cm 126.89 c4) 127.21 G) 126.96 (W
Piperidine
p(298.15 K)/g crnT3 nD(298.15 K) p(298.15 K)/Torr Tb(760 To&/K
0.85691 1.4505 30.49 379.45
0.85674 U) 1.4501 (3) 30.19 @) 379.55 (a
volume was about 50 cm3. The null manometer was a U-shaped glass tube. The inner diameter of the part of this tube used to read mercury levels was 14 to 15 mm. Both components were mixed in a glass vessel and degassed by repeating freezeand-melt procedure under vacuum. Degassed liquids were then transferred directly into the sample vessel without exposing them to air. Glass needle valves having Teflon needles were used wherever the sample liquids were allowed to pass through them. A Teflon-coated magnet was inserted in the sample vessel. It was driven by a submarine magnetic stirrer placed underneath the vessel and was proved effective in shortening the time required to attain vapour + liquid equilibrium. During the pressure measurements, the temperatures of the sample vessel and the null manometer immersed in a double thermostat were controlled within kO.01 K. The position of the thermostat was adjusted by lifting it with a laboratory jack. The total pressures of sample liquids at three different temperatures were read on mercury manometers by means of a cathetometer and a travelling microscope. Capillary corrections were applied to the manometer readings and the vapour pressures were then calculated by a conventional method using the local value of the acceleration of free fall and the density of mercury. The temperature of the thermostat was monitored by a Beckmann thermometer which was recalibrated against a calibrated standard thermometer. Its precision was +0.03 K. The vapour-pressure measurement was first done at 298.15 K, then the temperature of the thermostat was raised successively to 308.15 and 318.15 K, and the vapour pressure was measured at each temperature. The temperature was lowered again to 298.15 K and a duplicate measurement was made. Agreement with the first measurement was confirmed to complete one series of measurements. Equilibrium compositions were determined after the pressure measurements either by density or refractive index measurements. The densities of the sample mixtures were determined with a bicapillary pyknometer, the volume of which was about 30 cm3, and the refractive index measurements were made with a Pulfrich refractometer. Except for the piperidine-rich region where the density of the mixture goes through a flat maximum, the density measurement was preferred on account of
EXCESS
FUNCTIONS
OF
CHeOH
1127
+ CsHllN
better accuracy. The following smoothed values were used to convert the density p or the refractive index n, of a sample mixture at 298.15 K to the composition. P = bwl
+-M2(1 -xl))lwl/Pl)xl +w,IP,)(l --Xl> +x,(1-x,){ -5.O36-1.O25(2x,-l)+O.2771(2x1-1)2}
nn = nlxl+n~(~-x,)+x,(l-x,){0.1290+0.0578(2x,-1)+0.0174(2x,-1)2}. Subscripts 1 and 2 refer to methanol mass and x the mole fraction. EXCESS-ENTHALPY
and piperidine,
g cm-‘],
(1) (2)
respectively; M is the molar
MEASUREMENTS
An isothermal displacement calorimeter was constructed and used to measure the excess enthalpies at 298.15 K. Calibration of the calorimeter for exothermic mixing was done with dichloromethane -I- p-dioxane. The results(l’) agreed quite well with the values reported by Murukami and Benson.(“) To use this calorimeter in the present study, care had to be taken with the following problems. First, the O-ring packing of the Dewar vessel of the calorimeter must be completely resistent to highly corrosive amine-containing mixtures. A Viton E-60C O-ring was found to be suitable and used throughout the measurements. Secondly, the temperature rises on mixing are rather large for these liquids. To avoid the accumulation of errors on successive additions of the second component, the determination of composition must be done by calculations on a mass basis. Conversion of the volume added to the mass of the second component will lead to a non-negligible error inlthe composition. This situation has forced us to change the final design of the second-component feeding device of our calorimeter.(l’)
3. Results VAPOUR
In at to at
PRESSURES
table 2 are recorded the isothermal total pressures for x,MeOH + (1 -x&H, ,N 298.15, 308.15, and 3 18.15 K. Negative deviations from the Raoult’s law are found be fairly large. The errors in the vapour pressures are estimated to be f0.3 Torr 298.15 K and may be slightly larger at 308.15 and 318.15 K.?
EXCESS FUNCTIONS
From the above results, the excess Gibbs free energies GE were calculated by the Barker method.(13) Table 3 shows the values used for the molar volumeV$ and the second virial coefficient Bii of the pure liquids. Slightly different values of B,, are available for methanol(‘* ’ 6, and piperidine. u’) It was found that the calculated values of GE did not change significantly with the value assigned to Bii. There are no data on the cross virial coefficient Bij for this mixture. Therefore, 6, z = 2Bl 2 - B, 1Bz2 was assumed to be zero. t Throughout
this work
Torr
= (101.325/760)
kPa.
1128
K. NAKANISHI,
TABLE
2. Vapour
pressuresp,
H.
WADA,
excess enthalpies Crorr
AND
H.
TOUHARA
HE, and densities
= (101.325/760)
p for xlMeOH
+ (1 - xl)CSHllN
kPa)
p(298.15 ~(308.15 ~(318.15
Xl a K)/Torr K)/Torr K)/Torr
0 30.49 51.05 81.84
0.1196 32.83 56.00 92.1,
0.2058 35.7, 61.27 101.83
0.3002 39.6, 68.26 112.40
0.3868 42.90 73.7, 125.4,
0.4924 51.95 89.0, 149.0,
p(298.15 ~(308.15 ~(318.15
Xl a K).Torr K)/Torr K)/Torr
0.6754 75.43 127.54 208.2,
0.7829 94.1s 156.7, 252.59
0.8422 104.6* 173.6X 278.86
0.8815 110.84 183.56 293.9,
0.9393 120.15 197.6, 315.40
1 126.90 208.8, 333.2,
s”(298.15
K)/J
mol-l
-272.3 0.0309
-607.2 0.0670
-1130.5 0.1359
-1614.9 0.1989
-2032.1 0.2583
-2365.5 0.3091
%(298.15
K)/J
mol-l
-2687.2 0.3670
-2906.2 0.4170
-3047.8 0.4623
-3123.0 0.4940
-3186.3 0.5280
-3204.5 0.5546
G”(298.15
IQ/J
mol-l
-3205.8 0.5779
-3183.9 0.6062
-3107.6 0.6441
-2993.5 0.6775
-3020.0 0.6753
-2702.9 0.7317
%(298.15
K)/J
mole1
-2351.1 0.7791
- 0.8310 1902.4
-1361.6 0.8845
-766.0 0.9372
-310.9 0.9744
X1 p((298.15
K)/g
crnT3
0.1008 0.85773
0.2107 0.85700
0.3026 0.85642
0.3897 0.85502
0.5021 0.85182
Xl ~(298.15
K)/g
cme3
0.6949 0.83976
0.7974 0.82765
0.8571 0.81831
0.9054 0.80961
0.9547 0.79856
0.5992 63.06 107.80 177.41
0.5952 0.84729
4 Strictly speaking these x1 values are valid only at 298.15 K, since the sampling was done that temperature. Corrections to x1 for other temperatures were estimated to be less than 0.0001.
TABLE
3. Molar
volumes
Temperature T/K
Vt and second
virial
coefficients
Bii of the pure
MeOH V?/cm3
298.15 308.15 318.15
The Redlich-Kister
40.7 41.2 41.7
mol - 1
at
liquids
C&LN &/cm”
mol - 1
-1970 -1710 -1490
Vk/cm3
mol - 1
99.4 100.5 101.6
equation was used to express the composition
Bt,/cm3mol-1 -2116 -1910 -1740
dependence of GE :
GE/J mol- ’ = xlxz k A&, -xZ)nW1. n=l The results of iterative calculations indicate that the present results could be fitted satisfactorily by equation (3) with 4 constants, the values of which are given in table 4. The standard deviations o(p) of the vapour pressure are also included in the table. The excess enthalpies at 298.15 K are given in table 2. They were also fitted to an equation of the same type as equation (3) with HE replacing GE; the constants are given in table 4. The standard deviation o(HE) was 9.51 J mol-I. These two functions GE and HE were combined to calculate the excess entropies SE at 298.15 K.
EXCESS JWNCT’IONS OF CH,OH + C~HIIN
1129
TABLE 4. Coefficientsof the Redlich-Kister equation and smoothed values of excessfunctions for x = 0.5 (Torr = (101.325/760)kPa) T/K 298.15 GE
308.15 318.15
HE I/E
298.15
42
Al
-1239.5
-3336.2 -3100.3 -2764.2 - 12636.21
298.15 -5.036 Smoothed valuesfor x = 0.5 and T = 298.15 GE/Jmol - 1 HE/J mol-l -834.1 -3159.1
A
o(plTorr) 0.55 0.69
-1152.1 -1091.0
187.7 305.8 228.8
909.8 805.3 826.5
-3748.52
2038.28
2149.70
-1.025
0.2771
1.11
-
K.
TP/J mol - ’ -2325.0
P/cm3 mol - 1 -1.26
From the densities used to determine the equilibrium compositions, the excess volumes VE/cm3 mole1 were evaluated and fitted to a polynomial equation with three constants given in table 4. The standard deviation o(V’) was 0.014 cm3 mol-‘. Smoothed values of all the excess functions for x = 0.5 are also included in table 4. Figure 1 illustrates the dependence of GE, HE and TSE on x1. All the excess functions are fairly symmetrical with respect to x1, although minima are found at about
x1 = 0.55.
FIGURE 1. Excessfunctions of xlMeOH + (1 - xI)C5H11N at 298.15K. _---_
TSE;
-
XE:
-.-.-*.
HE.
- GE;
4. Discussion The excess functions of methanol
+ piperidine can be summarized as follows:
HE
VE< 0.
This is similar to the pattern observed for methanol + diethylamine or methanol + n-butylamine studied by us previously. (2) Comparing secondary amines, we find that TSE for MeOH + &HI ,N is less negative than that for MeOH + Et,NH, suggesting
1130
K. NAKANISHI,
H. WADA, AND I-I. TOUHARA
significant loss of internal degrees of freedom in E&NH molecule upon mixing. HE is also less negative for MeOH + C5H11N, in spite of the fact that the pK, of piperidine (11.22) is almost the same as that of Et,NH (10.98).(‘*) This implies that the effect of the basicity of the N-atom in amine molecules on the excess functions is of minor importance and that, in MeOH + C+H, ,N, there should be a packing effect which reduces the O-H---N hydrogen bonding energy. This interpretation may be supported by the fact that VE for MeOHf CSHllN is almost half of that for MeOH -I- Et,NH.@) Since the entropy effect mentioned above is predominant, GE for MeOH + C,H,,N is more negative than that for MeOH + E&NH. Several investigations including our unpublished work(lg’ have been made on an alcohol -I- pyridine. The excess functions for MeOH + C,H,N follow the same pattern as that given above. However, the absolute values of GE, HE, and TSE are smaller than those for MeOH f CSHIIN. The presence of n-electron interaction and NH-O hydrogen bonding, which might be responsible for this difference, can be proved to have only minor effects, and it may be concluded that the difference between the excess functions of MeOH -I- C,HSN and those of MeOH + C,H1,N can be ascribed to the weaker basicity (pK, = 5.23) of pyridine.(l*) The cooperation of Mr M. Kitamura in the construction of the vapour pressure apparatus is gratefully acknowledged. We also thank Professor N. Watanabe for stimulating discussions. REFERENCES
1. Nakanishi,K. ; Shirai, H. Bull. Chem. Sot. Japan 1970,43,1634. 2. Nakanishi,K.; Touhara,H.; Watanabe,N. Bull. Chem. Sot. Japan 1970,43,2671. 3. Riddick, J. A.; Bunger,W. B. Organic Solvents. Wiley: New York. 1970. 4. Manufacturing Chemists’Association,Data Sheets,table 23-2-1-(1.1000)-k, Thermodynamic Research Center,TexasA & M Univ., CollegeStation,Texas,1960. 5. Gibbard,H. F.; Creek,J. L. J. Chem. Eng. Data 1974,19, 308. 6. Polak, J.; Lu, B. C.-Y. J. Chem. Eng. Data 1972, 17,456. 7. Timmermans, J. Physico-Chemical Constants of Pure Organic Compounds. Elsevier:New York. 1965. 8. Osborn,A. G.; Douslin,D. R. J. Chem. Eng. Data 1968, 13, 534. 9. Wada,H. M. Eng. Thesis,Kyoto Univ., 1973. 10. Hermsen,R. W.; Prausnitz,J. M. Chem. Eng. Sci. 1963,18,485. 11. Touhara,H.; Ikeda, M.; Nakanishi,K.; Watanabe,N. J. Chem. Thermodynamics 1975,7,887. 12. Murakami, S.; Benson,G. C. J. Chem. l%ermodynamics 1969,1,559. 13. Barker,J. A. Aust. J. Chem. 1953,6,207. 14. Freydank,H.; Ratzsch,M. 2. Phys, Chem. (Leipzig) 1971,248,83. 15. Scott, D. W. J. Chem. Thermodynamics 1971, 3, 649. 16. Kudchadker,A. P.; Eubank,P. T. J. Chem. Eng. Data 1970, 15, 7. 17. Cabani,S.; Ceccanti,N. J. Chem. Thermodynamics 1973,5, 9. 18. Tamres,M.; Searles,S.; Leighly,E. M.; Mohrman,D. W. J. Amer. Chem. Sot. 1954, 76, 3783.
19, Nakanishi,K.; Ashitani,K.; Touhara,H. Unpublishedwork.