G. L. BERTRAND, J. W. LARSON, AND L. G. HEPLER
4194
Thermochemical Investigations of the Water-Triethylamine System by Gary L. Bertrand,l&John W. Larson,lb and Loren G. Heplerlb Department of Chemistry, University of Missouri at Rolla, Rolla, Missouri, the Department of Chemistry, Carnegie-Mellon University, Pittsburgh, Pennsylvania, and the Department of Chemistry, University of Louieville, Louisville,Kentucky (Received M a y $7, 1968)
Partial molar heats of solution of both water and triethylamine have been measured at 15’ over the whole range of composition in the water-triethylamine system. Integral heats of mixing have been calculated from these results and compared with earlier results obtained by differentmethods. Vapor pressure data have been used in calculating partial molar free energies and have been combined with partial molar heats to yield partial molar entropies, which indicate that triethylamine is a “structure maker” in dilute aqueous solution. Heats of solution of several nonelectrolytes have also been measured in water-triethylamine systems. Results are discussed in relation to heats of solution in water-ethanol and the effects of third components on critical mixing in water-triethylamine systems.
Introduction Phase equilibria in two-component liquid systems have long been of considerable practical and theoretical importance. Some of these two-component systems exhibit a lower critical solution temperature (lcst), in which case the liquids are completely miscible at low temperatures while they separate into two layers at high temperatures. The existence of an lcst has attracted considerable interest2 in terms of thermodynamic description of the phenomenon and also explanation in terms of molecular interactions. The water-triethylamine system, with lcst at 18-19” and X , (mole fraction of triethylamine) 0.08, has already been investigated in several ways. In order to extend the thermodynamic data for this system and to obtain information that can contribute to understanding the molecular interactions, we have determined both integral and differential heats of solution at 15”. Three-component liquid systems are also important in several respects. For example, many solvent extractions can be described in terms of a two-phase liquid system that consists of two principal components and a third dilute component. Also, many useful solvent systems consist of two “solvent” components and one (or more) “solute” components, often all in one liquid phase. I n addition to a general interest in systems of these kinds, we are specifically interested in the effects of third components on critical mixing in the watertriethylamine system. I n order to extend the earlier equilibrium investigations3 of effects of various solutes on critical mixing in this system, we have determined differential heats of solution of these solutes in watertriethylamine mixtures at 15’.
Experimental Section Three calorimeters were used in this investigation, one of which has been described p r e v i ~ u s l y except ,~ that a Leeds and Northrup Mueller G-2 bridge and an HS galvanometer were used with a nickel resistance The Journal of Physical Chemistry
thermometer. Also, the thermometer and manganin heater were contained in a glass spiral filled with mineral oil. This calorimeter, which held 950 ml of “solvent,” was used mostly for measurements in the water-rich region. Two smaller calorimeters were used for most of the measurements with triethylamine-rich solutions and for some with water-rich solutions. These smaller calorimeters, which will be described in detail later, were similar to the larger calorimeter described above. Each smaller calorimeter consisted of a dewar vessel that held 300 ml of solution. Temperature measurements were made with thermistors, in one system with a Mueller bridge and in the other with a transposed Maier bridge similar to that described previ~usly.~ All calorimetric measurements were made at 15.0 f 0.1’, in terms of the defined calorie. All measurements were made with freshly distilled water that was free of carbon dioxide. Triethylamine was refluxed over KOH and used shortly after it was vacuum distilled. Watertriethylamine solutions were prepared by volume, using calibrated glass ware. The ethanol used was U.S.I. absolute. Fisher sec-butyl alcohol was vacuum distilled. All other solutes were Fisher Certified chemicals.
Results and Discussion We have determined heats of solution of small quantities of pure water and pure triethylamine in much larger quantities of pure water, pure triethylamine, and watertriethylamine mixtures. Solute sample sizes (weighed to 0.1 mg) used in the large calorimeter (950 ml of “solvent”) ranged from 0.5 to 6 ml, while solute sample (1) (a) University of Missouri at Rolla, Rolla, Mo.; (b) University of Louisville, Louisville, Ky. (2) J. S. Rowlinson, “Liquids and Liquid Mixtures,” Butterworth and Co. Ltd., London, 1959. (3) B. J. Hales, G. L. Bertrand, and L. G. Hepler, J . Phys. Chem.,
(1966). (4) W. F. O’Hara, C. H. Wu, and L. G. Hepler, J . Chem. Educ., 38, 512 (1961). 70, 3970
THERMOCHEMICAL INVESTIGATIONS OF THE WATER-TRIETHYLAMINE SYSTEM sizes used in the small calorimeters (300 ml of "solvent") ranged from 0.1 to 2 ml. The molar heats of solution were found to be independent of solute sample size, except for solvent compositions approaching pure water. It is therefore appropriate to take most of our heats of solution to be differential or partial molar heats of solution. These differential heats of solution are given in Table I, where E, represents the molar differential heat of solution of water in a mixture of the indicated composition and Et indicates the molar differential heat of soluti.on of triethylamine. Compositions of solutions are given in terms of mole fraction of triethylamine, represented by Xt. Table I : Partial Molar Heats of Solution of Water and Triethylamine in Water-Triethylamine Mixtures at 15'
Xt
Et, ca,l/mol
0
0 0.00010 0.00030 0.0036 0.0071 0.0076 0.015 0.056 0.068 0.126 0.182 0.244 0.265 0.328 0.406 0.418 0.446 0.495 0.518 0.521 0.601 0.627 0.699 0.768 0.849 0,900 0,950 1.0 a
ZW,
cal/mol
-5400"
-8300' - 8920 -8370
- 9430" -2
- 22 - 126 - 135 - 194 - 284 -377 -431 -531 - 660 -661 -717 - 784 - 796 -815 - 982 - 1084 - 1200 - 1280 - 1310 - 1200 - 1000"
-6780 -2600
-2480
- 1890 - 1390
- 1000 -900 -666 -444
-347 -253 - 122 - 38 22 24 14 0
These values are calculated as described in the Discussion.
Results of several determinations of the heat of solution of triethylamine in water are given in Table 11. These results show that the heat of solution is strongly dependent on sample size, and it is therefore not justified t o take the observed molar heat t o be equivalent to the desired partial molar heat of solution. Instead, we treat these results as integral heats of mixing and use the method off Van Ness6 (graph of A H M / X t X wagainst X,) t o obtain the values indicated by footnote a in Table I. The value at X t = 1.0 was obtained by simple graphical extrapolation.
z, z,
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Table 11: Heats of Solution of Triethylamine in 16.57 Mol of Water Mol of triethylamine ( x 103)
cal
AHW'XwXt, koal/rnol of solution
0.691 0.692 1.371 3.450 3.507 6.972
4.3 4.2 10.2 28.4 29.1 61.4
-6.2 -6.1 -7.48 -8.24 -8.30 -8.81
Q,
We have tested the internal consistency of our data by means of the equation
Using E , data with this equation, we have calculated Lt values that agree well with the experimental Et values. Similarly, the experimental Et values lead to L, values in good agreement with the experimental results. We can see from eq 1 and the E, data that L, values must pass through a maximum for solutions very dilute in triethylamine, but the E, values in this region are too small for us t o measure accurately. We have used our values to calculate integral heats of mixing to form 1 mol of solution by means of the equation
AH^
=
x,Z,
+ x,Z,
Results of these calculations are shown in Figure 1, where we also show results obtained by other investigators. We estimate a maximum uncertainty of 10 cal/mol in our A H M results. The integral heats of mixing measured by Copp and Everetts agree with our results only for solutions having a small Xt. As is common for integral heats of mixing, their results are inadequate for calculating differential heats to compare with our values cited in Table I. The integral heats of mixing calculated by Kohler' by differentiation of his vapor pressure data determined at several temperatures are in even poorer agreement with our results. We believe that the disagreement may be attributed almost entirely to the large uncertainty associated with Kohler's heats, although there is no reason to question the free energies of mixing reported by Kohler.7 MatiZen and Kuskova* have made calorimetric measurements leading to values for several solutions having compositions up to X , g 0.4. Their experimental (5) H.C.Van Ness, "Classical Thermodynamics of Non-Electrolyte Solutions," The Macmillan Co., New York, N. Y.,1964. (6) J. L.Copp and D. H. Everett, Discussions Faraday SOC.,15, 174 (1953). (7) F. Kohler, Monatsh. Chem., 82, 913 (1951). (8) E. V. Matizen and N. V. Kuskova, Zh. Fiz. Khim., 34, 2223 (1960).
Volume 72,Number 12 November 1968
G. L. BERTRAND, J. W. LARSON, AND L. G. HEPLER
4196
-
70C
0 0 0
-801
0
-SO(
-400
-300
-20.0
-100
d.1
0:2
0!3
0.4
I
1
0.5
0.6
0.7
Oh
I
0.9
xt
Figure 1. Heats of mixing of the water-triethylamine system: , this work; A, Copp and Everett; 0, Kohler; B, Matizen and Kuskova (experimental); 0, Matizen and Kuskova (from eq 1).
values are in satisfactory agreement with our values in Table I, except for solutions with very small Xt for which neither their nor our results are highly accurate. Their integral heats of mixing calculated from eq 2 are in satisfactory agreement with our values. They* also used their E, values with eq 1 t o calculate some Et values, which in turn were used with eq 2 in calculating heats of mixing. These latter heats are not in good agreement with our values, nor are their calculated Et values in good agreement with our measured Et values. I n summary, we believe that our and AHM values, which are supported by the experimental results (but not all the derived quantities) of Matieen and Kuskova,8 are the best thermal data presently available for the water-triethylamine system. From the data of Kohler’ we have calculated partial molar excess free energies for both water and triethylamine at 10 and 18”, neglecting corrections for nonideality of the vapors, and interpolated to find values at 15”. We have combined these quantities with our E data to obtain excess partial molar entropies represented by SE. Results of these calculations are given in Table 111. The entropies show that triethylamine is a “structure maker’’ in dilute aqueous solutions, which is relevant to results reported and discussed in the remainder of this paper.
aE
The Journal of Physical Chemistry
Table I11
Xt
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0
oal/mol
1173 782 560 407 285 175 114 54 11 0
oal/mol
0 40 107 179 26 1 357 472 613 786 1005
mol-’
-11.3 -7.1 -4.55 -2.94 -2.22 -1.06 -0.52 -0.15 0.05 0
mol-’
0 -0.75 -1.46 -2.28 -3.12 -3.98 -4.91 -5.9 -7.1 -8.1
We have measured heats of solution of very small quantities of eight nonelectrolyte solutes in much larger quantities of water-triethylamine mixtures. The results, which closely approximate differential heats of solution of solute at infinite dilution, are given in Table IV. There are certain similarities in the heats of solution of alcohols in triethylamine-water mixtures (Table IV) and heats of solution in various alcohol-water mix-
THERMOCHEMICAL INVESTICATIONS OF THE WATER-TRIETHYLAMINE SYSTEM
4197
Table IV : Heats of Solution (kilocalorie per mole of solute a t infinite dilution) in Triethylamine-Water a t 15’ 0
Ethanol Propanol 1-Butanol 2-Butanol see-Butyl alcohol t-Butyl alcohol Benzene Chloroform
-
Xt
,-
0.024
0.055
-2.74 -2.86 -2.67 -2.80 -3.69
-1.87 -0.85 1.25 1.01
-1.55 -0.25 1.75
-3.48
-1.12
...
17.5 20
...
...
0.079
... -0.24 1.10 1.55
0.114
...
...
0.46
...
0.54 0.91 0.99
0.26
...
0.59
...
5.06 4.00
...
11.0 11.7
1
t
.
t u r e ~ . ~ ,For ’ ~ example, heats of solution of all solutes are more endothermic in highly aqueous mixtures than in pure water and commonly go through a maximum at X , > 0.7. The height of the maximum is greater for large solute molecules (butanols) than for smaller solute molecules (ethanol) in water-triethylamine mixtures, just as previously observedg for waterethanol mixtures. Since the entropy data (Table 111) indicate that triethylamine is a “structure maker” in dilute aqueous solutions just as similar data indicate that ethanolg~’Oand other alcoholsg are “structure makers” in dilute water-alcohol mixtures, the qualitative explanation previously givens for water-alcohol mixtures is applicable to the water-triethylamine mixtures under consideration here. Qualitative effects of third components on mutal solubilities of liquids have long been known for watertriethylamine and other systems. Third components that are very soluble in both components of the “solvent” commonly increase their mutual solubility. This increase in mutual solubility corresponds to raising the mutual solubility curve ( A T positive) for the water-triethylamine system. On the other hand, solutes that are very soluble in either water or triethylamine but only slightly soluble in the other liquid commonly lower the mutual solubility curve ( A T negative). Results of previous investigations3 have been reported in terms of a and b values taken from A T = aXa bXS2,in which XBrepresents the mole fraction of added third component. The values of a and b were found to depend on the particular solute and upon X t . Of particular interest are those solutes for which a values change sign with increasing X,. These solutes
+
0.151
-1.25 -0.52
... ... ...
...
0.244
-1.16 -1.08
0.381
-1.29
...
...
-1.04
-0.16 -0.22
...
0.53
...
...
...
0.35 1.06
...
0.500
-1.21 -1.18 -1.26 -1.16 -0.97
0.627 I
.
.
... ...
... ...
0.34
0.34
...
0.37
-3.15
...
0.742
-1.16 I
.
.
-1.22
... -1.02 0.44
... -3.40
1,000
-0.93 -1.03 -1.02 -1.03 -0.54 1.05 0.24 -3.42
lower the mutual solubility curve (both A T and a negative) for mixtures having very small X t and raise the mutual solubility curve (both A T and a positive) for mixtures with larger Xt. Thus at some intermediate X t and corresponding temperature, these solutes have no effect on the mutual solubility of water and triethylamine. The alcohols (1-butanol, 2-butanol, secbutyl alcohol, and 1-propanol) that exhibit this behavior are those that have sharp peaks in the A H of solution us. X, curves that can be constructed from the data in Table IV, while those alcohols (t-butyl alcohol and ethanol) for which A T and a are positive for all X , are those whose A H os. X,curves are quite flat on the high X t side. We also note that the trend in values of a for water-triethylamine mixtures with Xt near that of the lcst is paralleled by the shift in maxima of the AH of solution curves to larger X t . The heat of solution of benzene in triethylamine is quite small, as expected for liquids that do not interact strongly. On the other hand, the heat of solution of chloroform in triethylamine is exothermic as expected on the basis of the known formation of a stable compound in the chloroform-triethylamine system. l1
Acknowledgment. We are grateful to the National Science Foundation for support of part of this research. (9) E. M. Arnett. W. J. Bentrude. J. J. Burke. and P. McC. Duegleby, J . Amer. Chem. Soc., 87, 1541 (1965); E. h.Arnett and D. 6. McKelvey, ibid., 87, 1393 (1965). (10) G. L. Bertrand, F. J. Millero, C. H. Wu, and L. G. Hepler,
J . Phys. Chem., 70, 699 (1966). (11) G. W. Stapleton, M . Bellay, C. A. Wulff, and L. G. Hepler, J . Chem. Eng. Data, 11, 96 (1966).
Volume 79, Number 12 November 1968
