F. KOHLER,E. LIEBBRMANN, G. MIJ[CSCW, AND C. KAME
2764
ioiis and simple NOs- anions. This conclusion is in J ~system. J ~ agreement with previous W O ~ ~ ~ -on~ this
Acknowledgment. 5. N. wishes to acknowledge the award of a 'Rotary Fellowship.
On the Thermodynamics of the Acetic Acid-Triethylamine System1 by Friedrich Kohler,* E. Liebermann, G. Miksch, and Christine Kainz Imtitute of Physical Chemistry, University of Vienna, Vienna, Austria
(Received March 80, 1978)
The system acetic acid-triethylamine is one of the rare examples where phase separation occurs at negative values of the excess Gibbs free energy of mixing AGE. The vapor pressure curve exhibits a negative azeotrope besides the miscibility gap. Activity coefficientsare calculated for 20" from the melting curve on the acetic acid wide, from the vapor pressure of the negative azeotrope, from the concentrations along the coexistence curve, and from the vapor pressures between the miscibility gap and pure triethylamine. Orientative values for the heat of mixing are given. All properties indicate the existence of a very stable aggregate consisting of three molecules of acetic acid plus one molecule of triethylamine in the liquid state. Determination of tho excess volumes shows an exceptionally large contraction around that mole ratio; viscosities have a very sharp maximum. The results are interpreted by a largo attractive interaction between the polar 1:1. complex (aeot!c acid-triethylamine) and the acetic acid dimer, which causes the liquid mixture of equimolar concentration to split into a phase rich in the 3: 1 aggregate and a phase rich in triethylamine.
Introduction The peculiar properties of the system acetic acidtriethylamine have been emphasized by several aut h o r ~ . ~ It - ~has a negative azeotrope of very low vapor pressure a t z1 (mole fraction of acetic acid) = 0.75, which stays ai; this composition a t least until boiling a t atmospheric pressure. For smaller mole fractions of acetic acid, there is a region of phase separation. Some measurements of total pressures have been carried but a meaningful thermodynamic evaluation would have needed a proper consideration of the dimerization of acetic acid in the vapor phase, In order to obtain the thermodynamic properties of the system, we have measured vapor pressures on the triethylamine side and concentrations of coexisting phases. On the acetic acid side, we have determined the melting curve and the vapor pressure curve of the negative azeotrope. The vapor pressure measurements are evaluated on the assumption that the vapor COIIsista of triethyhmine molecules, acetic acid monomers, and acetic acid dimers. Triethylamine-acetic acid complexes are neglected in the vapor phase. This certainly does not introduce any error on the triethylamine side, but it) inay affect 1,lze evaluation from the vapor pressure of the azeotrope. However, the conclusions drawn from the vapor pressure curve of the azeotrope fit into the general picture. In previous discussions on acetic acid mixtures5z6 attention has been called for the apparent existence of a strong interaction between a polar molecule and the The Journal
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Physical Chemistry, Vol. 76, No.
19, 197%
distortable hydrogen bonds of the acetic acid dimer. This interaction is accompanied by a large negative change of free energy and by a large volume contraction. Assuming a highly polar form of the 1 : 1 complex between acetic acid and triethylamine, this interaction can explain the preference for aggregates between acetic acid and triethylamine in the mole ratio 3: 1. Volumetric measurements show indeed an extremely large volume contraction, which even increases with temperature. This is contrary Lo the behavior of hydrogen-bonded complexes, like EtyN. 3Hd3, where the complexes break up with increasing tornperature, leading t o a decrease of the volume contraction. It has not been possible to obtain crystals around the mole ratio 3 : 4. ;instead the formation of a glass occurred. The viscosities have been measured and show a very sharp maximum near this mole ratio.
Experimental Section and Results Acetic acid (Pro Analysi) was fractionally disl.illed (1) Presented in part at the 161st National Meeting of the American Chemical Society, Lo8 Angeles, Calif., March 1971, (2) 11. S. van Klooster and W. A. Douglas, 6.Phgs. Cham., 49, 67 (1945). (3) J. Hollo, T.Lengyel, and H. M. Uaonyi, Period. Polglech. Chem. Eng., 4, 172 (1960). (4) V. F. Plekhotkin and N. P.Markuzin, E'iz. Khim. SvoisEva RastPO TO^, 12 (1964); A. V. Storonkin, N. P. Markupin, and V. F. Plekhotkin, {bid., 19 (1964). (5) H.E. Affsprung, G. H. Findenegg, and F. Kohler, J ,Chem. Xoc. A , 1364 (1968). (6) F. Kohler, Monalsh. Ckem., 100, 1151 (1969).
T H ~ R M ~ ~ YOFNTHE ~M ACETIC I C ~ACID-TRIETHYLAMINE SYSTEM
2165
in a 25-plate column (reflux ratio 20: 1) after addition stant boiling point. The mixture was degassed and of some acetic acid anhydride. It was stored in the its vapor pressure measured statically. These results dark, the wrrpor plzase being in contact with P2Q,. The are shown in Table 111. J ) ) 1.3716, the density ( ~ 2 0 ) refractive index ( ' Y L ~ ~was was I. 04952 g/mn3jarid the melting point was 16.53*. ~ ~ ~ ( ~ ' ~ ~ was ~ ~shaken u ~ t with j ~solid KCIB, y ~ Table~ILI: Vapor ~ Pressure ~ atn the Composition ~ of bhe Negative Azeotrope (Mole Fraction Acetic Acid 0.75) fractionally d i ~ t ~twice l ~ ~over d sodium wire, and stored in the System Acetic Acid-Triethylamine in the dark aver solid XOH. The refractive index (+OD) t, "G 40 50 60 70 nsity ( p 2 0 ) was 0.72729 g/cm3. 80 158.0 P, m m H g 1.99 3.84 7.35 13.28 23.30 742.5 ureh were determined as for the system acetin carbon tetrachloride.? The i~verain Tatdc I. The vapor pressure curves The melting curve on the acetic acid side was deteron the t r ~ e t h ~ ~side ~ a of ~ the ~ ~ miscibility ne gap lie below mined with a method described elsewhere.8 The rethe vapor presmre curve of an ideal mixture a t 20 and sults of the measurements are given in Table IV. 30" m d slightly above t h a t at 40*.
-
Table I V : Melting Point Lowering of Acetic Acid in the System Acetic Acid-Triethylaminea
Table X : Total Vapor Pressure P of Solutions of ,4cetic Acid in Triei,hylamine"
-
P , rnrn Hg51
200
;30°
400
0.0000 0.0223 0 I0353 0 "0612 0.0786 Two -phase region I * 0000
53.36 51.95 51.30 49 62 48.75
86.12 84.23 83.26 81.41 79,82 74.57
132.12 130.06 129.02 126.31 124.61 117.42
0
51
~
45,29
35.29
11.76
= mole fraction of acetic acid.
The consolute curve was determined by weighing the mixture in a tube closed by a ground stopper, and with an insert which allowed vigorous shaking without wetting of the ground stopper, The tube was placed into a thermostat and both the disappearance of any turbidity when slowly raising the temperature and the appearance of the first turbidity when slowly Iowering the temperature were noted. The measured points are listed in Table 13;. A mixture of the mole ratio 3: 1 of acetic acid was prepared and distilled. The mixture showed a conI
---
,
Table I1 : Concentrations of Phase Separation in the System Acetic Acid-TriethylaInine between 20 and 450a
1, o c
I 1
1.0000 0.9849 0.9751 0.9583 0.9484 $1
=
16.53 15.51 14.69 12.85 11.42
a1
0.9249 0.8880 0.8525 0.8391
t , "C
6.48 -2.99 -14.44 -19.00
mole fraction acetic acid.
The pycnometric determination of densities has been d e ~ c r i b e d . ~For the tabulation of the results see the microfilm edition.Io Viscosities were determined in Ostwald vrscosimeters calibrated with water. Results arc, shown in Figure 1 (the tabulation is given in the microfilm edition).l0 Thermodynamic Evaluation The melting curve furnishes the activity coefficient2 fi of acetic acid at a reference temperature Yrof according to the equation I;" AT Infl = --
RTT,'
ACplrAT2 -
SRTT,
I n eq 1, T is the melting temperature of the mixture, I', that of pure acetic acid, and AT = ,'Y - T . The heat of melting L" of acetic acid can be checked by (7) G. Miksch, F. Ratlrovics, and E'. Kohler, J . Chern. Therrnodyn.,
t , "C
21.5 24.4 25.1 33.1 38.6 39.1 44.3 a
ZI =
1, 257 (1969).
Z .-.1 -
0.3375
0.5919
0 1467 ~
0.1541 0.1553 0.1568
mole fraction of acetic acid.
0.5907 0.5874 0.5865 0.5836
(8) E. Liebermam and I?. Kohler, Monatsh. Chem., 99, 2514 (1968); R. J. Munn snd F. Kohler, (bid., 91, 381 (1960). (9) 6. N.Findenegg and F. Kohler, Trarm. Faraday Soc., 6 3 , 870 (1967). (10) Tabulation of the densities and viscosities will appear following them pages in the microfilm edition of this volume of the journal. Single copies may be obtained from the Business Operations Office, Books and Journals Division, American Chemical Society, 1155 Sixteenth St., N.W., Washington, D. C. 20030, by referring t o code number JPC-72-2764. Remit check or money order for $3.00 for photocopy or $2.00 for microfiche.
The Journal of Physical Chem.istry, Val. 76, N o . 15,1971
F. KOHLER, E. LIEBERMANN, G. MIBSCIZ,AND C. KAXNZ
30
"lP)
IO
-
"7
Figure 1. Viscosity 9 of the system acetic acid (1)-triethylamine at 25 and 40" for mole fractions on the acetic acid side of the miscibility gap.
extrapolating the quantity RTT, (1. - XI)/ATto XI = 1. Such a plot is linear for small values of 1 - XI if In f l can be approximated by a parabola (Figure 2 ) ; it leads t o L" = 2780 cal mol-' in good agreement with the literalure11~*2 (2757 and 2803 cal mol-l, respectively). The difference in the heat capacities at melting AC," for acetic acid is taken as 5.4 cal "K-l mol-.' by a rough extrapolation of C p values of both solid and liquid phases.12 The partial molar heat of mixing AH1 is neglected. This will be discussed below. The vapor pressure P of the azeotrope gives directly the activity coeEcients of both components, if the vapor can be assumed to contain an ideal mixture of triethylamine, acetic acid monomeres, and acetic acid dimers
Infi = In
.P Po2
4K,P(2 - Xl) - s + In 1 + 2K,P(2 X1)2
with S = dI $ 4KpPx1(2- xl), K , being the dimerization constant of acetic acid in thc gas phase, and Pal and P,z denoting the vapor pressures of the pure components. The branch of the total pressure curve on the trlethylamine side of the miscibility gap can be evaluated by a differential equation derived from the GibbsDuhem equation, which is modified to allow for dimerization of acetic acid in the vapor7
dYl dP
I
--+(l-X,l
-
Figure 2. The function IZTTm(l a ) / A T US. (1 - SI); extrapolation t o x1 = 1 gives the heat of melting of acetic acid.
Kere y1 is the (formal) mole fraction of acetic acid in the gas phase and S is defined (in the general case) by S = dl 4KpPy,(2 - yl). The numerical integration of eq 3 has to start a t the lowest total pressure, ie., the pressure over the two-phase region. The corresponding value of yl", which is needed as initial value for the integration procedure, is not known. Therefore, integration has been carried out, for various values of yl". It is clear that the choice of yI* influences strongly the position of the In fl curve on the triethylamine side, but is relatively unimportant for the In f 2 curve, which is near zero anyway, Therefore, the value of In ,f~on the acetic acid side of the twophase region is mainly determined by the measurements of the consolute curve. I n order to set the various pieces of information together, first the In fi curve on the acetic acid side bas been constructed from the points at the consolute curve (XI = 0.59 a t 20") and at the azeotrope (21 = 0.75) and from the known slope in the interval 0.84 5 2 1 5 I , which can be obtained from the fl values determined by the melting curve by using the Gibbs-Duhern rclation. The construction of the lnf2 curve can be checked by recalculating the In $1 curve, which has to fit the points of the melting curve and the value at the azeotrope. Extending the Gibbs-Duhem integration of the In ,fl curve now to the two-phase region, the correct, value of yl* can be deduced. The behavior of the activity coefficients in the total concentration interval tiorl is shown in Figure 3; the probable ~ ~ o n ~ ~ ~ ~ofu athe curves within the interval of instability of a homogeneous liquid is also indicated. The total Gibbs free energy of mixing (excess plus ideal) is given in Figure
+
(11) J. Meyer, 2.Phys. Chem., 72, 225 (1910). (12) G . S. Parks and K. K. Kelley, J. Amer. Chem.
(19%).
The Journal of Phgssical Chemistry, Vol. 76, No. 10, 1078
SOC., 47,
2089
TNERMODYNAMICB OF THE ACETICACID-TRIETHYLAMINE SYSTEM
2167
-20011
t
““I---1000
A G V ** ~,
-8000
-10000
AH^/^^^^ Figure 3. The logarithm of the activity coefficients of the system acetic acid (1)-triethylamine (2) at 20”. The calculated functions in the range of instability are indicated by dashed lines.
-12000
cal/mol
Figure 5 . The excess Gibbs free energy of mixing (full curve) and the heat of mixing (dashed curve) divided by the product of mole fractions of acetic acid (1)-triethylamine at 20’.
the melting curve which have the largest temperature difference, a positive AN1 value in eq 1 is necessary for a smooth continuation of the In fl curve to the points deduced by the azeotropic pressure. Figure 6 shows the values for the volume change of mixing divided by the product of the mole fractions. The abnormally large contractions are remarkable.
Discussion AG%AG i’
I
c al/m D I
Figure 4. Ideal and excess Gibbs free energy of mixing of the system acetic acid (1)-triethylamine (2) at 20’. The dashed parts of the curves refer to the range of instability.
4 to illustra,te the phase separation. The excess, divided by the product of the mole fractions, is given in Figure 5. Values for the heat of mixing divided by the product of the mole fractions are also included in Figure 5 . These values are based on orientative calorimetric measurements at both ends of the concentration interval, on the temperature dependence of the azeotropic pressure, and on the temperature dependence of the total pressure near the consolute curve. They may be in terror by i 1 O O j , . At the acetic acid side, the curve for AH 21s. 21seerrts to be curved in such a way that the partial heat of mixing of acetic acid AH1 has a positive portion for high concentrations of acetic acid. This is in agreement with the fact that for the two points of
Figure 5 shows clearly that the formation of the I : 1 complex between triethylamine and acetic acid, which occurs on the triethylamine side, is accompanied by a negative AG and by an about equally negative TAS contribution, such that AH of the complex formation is about twice as negative as AG. I n contrast t o this, the formation of the aggregate triethylamine-3 acetic acid, which occurs on the acetic acid side, is accompanied by a strongly negative AG contribution and a comparatively small, though negative, TAX term. This can only be interpreted by a relatively indefinite orientation of the constituent molecules in the triethylamine4 acetic acid aggregate. This conclusion is confirmed by the measurements of the volume change of mixing. Let us compare the situation with that in pure acetic acid. The dimerization of acetic acid occurs without a volume change.13 Apparently, the closer proximity of the two monomeric groups is compensated by the larger rotational volume of the rigid dimer. On the other hand, there is good (13) T. A. Litovitl; and E. Carnevale, J. Aeoust. Soe. Amer., 30, 134 (1968).
The Journal of Physical Chemistry, Vol. 76, N o . 19, 1972
E. KOHLER, E. LIEBERMANN, a. hIIKSCH,
2768
-5D*1
10'
t
AND
c. I(A1NZ
i
-40
I
8
#,I
I
-.+x1 Figure 6. The volume change of mixing divided by the product oi mole fractions of acetic acid (1)-triethylamine at 201 and 40'.
indication for a volume contraction in the monomerdimer interact,ionJ6which is thought to be relatively unspecific with respect to orientation. Similarly, the well-defined complex triethylamine-3 water shows a much smaller ~ o n t r a c t i o n ~ than * , ~ ~ the aggregate triethylamine-3 acetic acid. Furthermore, the rigid complex breaks up with increasing temperature, which is not the case for the aggregate triethylamine-3 acetic acid, as shown by a comparison of the expansion coefficients of the two systems (Figure 7). The aggregate triethylamine-3 acetic acid, which is very stable but orientationally not well defined, is
The Journal
of
Physical Chemistry, VoE. 76, No. 19, 1878
3 x 8
Figure 7. Mean expansion coefficient a : acetic acid (1)-triethylamine, temperature range 20-40' (full curve); water (1)-triethylamine, temperature range 0-13" (dashed curve).
thought to be caused by a strong inductive interaction between the polar 1:1 complex and the distortable hydrogen bonds of the acetic acid dimer.
Acknowledgment. The authors wish to thank i\/liss Brucher for carrying out the orientative measurements of the heat of mixing and the Institute of Statistics of our University for computer time. (14) E'. Kohler, H.Arnold, and R. J. Munn, Monatsh. Chem., 92, 876 (1961); F.Kohler, ibid., 82, 913 (1951). (15) R. Schano, Doctoral Dissertation, University of Vienna, 1969.
