2524
J. Phys. Chem. 1901, 85, 2524-2529
energy of activation as defined by Eyring's equation q = (hNA/v)eAC'/RT
(4)
where u denotes the molar volume, is plotted against M. Such a graph, for several paraffins, tertiary amines, and carboxylic acids, is shown in Figure 8. Table I1 shows the relative molecular masses for various complexes M in comparison with values MARC obtained empirically via Figure 8. The following conclusions can be drawn. For alkanoic acids and tertiary amines, the value M- obtained from viscosities at ?&id = 0.75 corresponds rather well to that expected for a 3:l aggregate. The enhancement M ~ J Mobserved for aggregates with di-n-butyl- and nbutylamine appears to be too large to be an artifact of our empirical method and is perhaps caused by additional attachment of ionized groups. For trifluoroacetic acid, the 3 1 aggregate seems to be rather weak. On the other hand, trifluoroacetic acid forms a very strong 1:l complex with triethylamine, in the form of a zwitterion,as we will discuss in the following paper,l' and these zwitterions seem to be (17) Kohler, F.; Gopal, R.; GWze, G.; Atropa, H.; Demiriz, M. A,; Liebermann, E.; Wilhelm, E.; Rakovics, F.; Palagyi, B. J.Phys. Chem., following paper in this issue. See also ref 7 rind 5.
partly aggregated. Still higher aggregation can be found, at equimolar composition, with di-n-butylamine and especially n-butylamine, where the cluster corresponds to ca. four 1:l complexes. Acknowledgment. The cooperation between the Ruhr-Universitat Bochum and Veszprgmi Vegyipari Egyetem has been supported by the Deutsche Forschungsgemeinschaft and the Office for Cultural Relations of the Hungarian People's Republic, for which we are thankful. Supplementary Material Available: Tables containing the following: (i) melting points for the systems formic acid (1)+ triethylamine (2) (for 0 < x 2 C 0.98), propionic acid (1) + triethylamine (2) (for 0 C x2 C 0.75); (ii) dynamic viscosities of the systems formic acid + triethylamine, trifluoroacetic acid + triethylamine, propionic acid triethylamine, and trimethylacetic acid + triethylamine for various mole fractions a t 293 and 313 K,and for the systems propionic acid + tri-n-butylamine, propionic acid + di-n-butylamine, propionic acid + n-butylamine, and trimethylacetic acid + tri-n-butylamine for various mole fractions and from 293 to 323 K (10 pages). Ordering information is available on any current masthead page.
+
Molecular Interactions in Mixtures of Carboxylic Acids wlth Amines. 2. Volumetric, Conductimetric, and NMR Properties Frledrlch Kohler,' Ram Gopal,+ G. Gotze, H. Atrops, M. A. Demlrlr, Lehrstuhi fur Thermodynamik, Ruhr-Universitat Bochum, Postfach 10 2 I 48, 0-4630 Bochum, German Federal Republic
E. Llebermann, Emmerlch Wllhelm, Institut fur Physikaiische Chemie, Universitat Wien, A- 1090 Wien, Austria
F. Ratkovics, and B. Palagyl Vegylparl Egyetem, Fizikai K6miai Tanszgke, H-820 1 Veszprgm, Pf. 28, Hungary (Received: January 27, 198 1)
In this series of articles, we attempt to present a unified treatment of systems of carboxylic acid + amine. In part 1,experimental results on melting curves and viscosities were reported. Here, we present excess volumes for formic, trifluoroacetic, propionic, and trimethylacetic (pivalic) acids + triethylamine, as well as their temperature dependence, and excess volumes for propionic acid + n-butylamine and propionic acid + di-nbutylamine. Electrolytic conductivitieshave been measured for formic, trifluoroacetic, acetic, and propionic acids + triethylamine,propionic acid + tri-n-butylamine,propionic acid + di-n-butylamine,and propionic acid + n-butylamine. In addition, chemical shifts for trifluoroacetic, formic,acetic, and propionic acids + triethylamine, propionic acid + tri-n-propylamine,propionic acid + tri-n-butylamine, and propionic acid + n-propylamine are represented.
Introduction In the first paper of this series,l we have summarized the peculiar properties of mixtures of carboxylic acids with amines in terms of molecular interaction and complex formation. Without going into details, we recapitulate briefly the basis of our interpretation. For most mixtures we assume (1)predominance of cyclic dimerization of the carboxylic acid, (2) formation of a strongly polar 1:l complex between carboxylic acid and amine, and (3) a very attractive interaction between this polar complex and the t C-1087 Sector A, Mahangar, Lucknow-226006, U.P., India. 0022-3654/81/2085-2524$01.25/0
easily distortable ring (containing two hydrogen bonds) of the cyclic dimer. This gives rise to the formation of orientationally ill-defined 3:l aggregates with extreme influence on thermodynamic and structural properties. The only acid where these 3:l aggregates seem to be less important relative to the 1:l complex is trifluoroacetic acid, where the 1:l complex is a zwitterion. With formic acid, where open dimers and higher chain associates play a dominant role in the pure acid, the 3:l aggregate seems (1) Kohler, F.;Atrops, H.; Kalali, H.; Liebermann, E.; Wilhelm, E.; Ratkomcs, F.; Salamon, T. J.Phys. Chern., preceding paper in this issue.
0 1981 American Chemical Society
Mixtures of Carboxylic Acids with Amines
The Journal of Physical Chemistty, Vol. 85, No. 17, 1981 2525
to be partially displaced by short chains terminated by an amine molecule. In the case of alkanoic acids with bulkier groups, e.g., trimethylacetic acid, the dominance of the 3:l aggregates, though still seen in some properties, becomes less decisive for the thermodynamicbehavior. At the same time, a 2:l complex increases in importance. In this paper we will substantiate our ideas by presenting volumetric, conductimetric, and NMR properties, thus discussing also more elaborately questions concerning ion formation. In this connection, special reference should be given to a recent paper of Huyskens et aL2 Experimental Section References for purification and storage of carboxylic acids and amines have been given before,l together with some properties of the pure components (see Table I of the preceding paper1). The excess volumes were computed from densities, measured with a vibrating tube densimeter DMA 02D (Anton Paar, Graz, Austria); some check measurements were done with 15-cm3 double stem pycnometer^.^ All mixtures were prepared by weight and corrected for the vapor phase. Electrolytic conductivities were measured with a conductometer of Philips (PW 9501). NMR chemical shifts were measured with a Varian A-60A spectrometer with internal temperature regulation (40"C);tetramethylsilane was used as internal standard. Some of the propionic acid systems were measured with a 80-MHz TESLA BS 487C at -25 "C,with hexamethyldisiloxane as external standard. Some of the scattering of Figure 11 might be caused by these experimental differences, but the more critical fact is the purity of the substances.
Results The direct experimental values of excess volumes and electrolytic conductivities are tabulated and are available as supplementary material (see paragraph at end of text regarding supplementary material). The NMR chemical shifts are presented in Figures 8-12. Discussion Volumetric Measurements. In Figure 1,the molar excess volumes uE at 293.15 K are shown for the systems formic, propionic, trifluoroacetic, and trimethylacetic acids triethylamine (for the last system, the density of trimethylacetic acid was extrapolated to 293.15 K from eight measurements between 308 and 338 K). We present data as uE in Figure 1A and in the form uE/(x,x2) in Figure 1B. The first plot would give straight lines for complete complexation, whereas the second shows more directly the magnitude of the contraction in an average molecular encounter. In Figure lB, the curve for acetic acid + triethylamines is shown for comparison. It is remarkable that the results for uE of acetic acid + triethylamine and for propionic acid triethylamine coincide almost completely, though the former system shows phase separation, whereas the latter does not. Values for uE of trimethylacetic acid triethylamine are very similar over the whole composition range, but somewhat less negative
+
+
+
(2) Huyskens, P.;Felix, N.; Janssens, A.; Van der Broeck, F.; Kapuku, F.J. Phya. Chern. 1980,84, 1387. (3) Siddiai. M. A.: G6tze.. G.:. Kohler, F. Ber. Bunsenges. Phys. Chern. 1980; 84,529.' (4) Findenegg, G. H.; Kohler, F. Trans. Faraday SOC.1967, 63,870. (5) Kohler, F.; Liebermann, E.; Miksch, G.; Kainz, C. J. Phys. Chem. 1972, 76, 2764.
-15
\ CF, \
i
'r
\
1 0
I
1
0.2
0.4
iI
1
06
0.8
1.0
x2
-io
I
I
I
I
/I
i Figure 1. (A) Volume change of mixing vE of the systems formic (H), trifluoroacetic (CF,), propionic (CH3CH,), and trimethylacetic ((CH&C) triethylamine (x,) at 293.15 K vs. mole fraction x p . Apacids proximate curves are drawn to illustrate complexation, points denote As in A, but for the quantlty $/(x1x2). Curves experimentalresults. (6) were calculated according to eq 1.
+
than for the other systems. These three systems show clearly the dominance of the 3:l aggregation, for which contraction is extremely strong. The shift of the intersection point in Figure 1A from x 2 = 0.27 for propionic acid + triethylamine to x 2 = 0.29 for trimethylacetic acid + triethylamine might indicate a slightly enhanced contribution of the 2 1 complex in the latter system. Formic acid + triethylamine has a much smaller contraction at high mole fractions of acid, but it increases until phase separation. Trifluoroacetic acid triethylamine has the strongest contraction. Here the dominance of the 1:l complex is emphasized, but some changes of the aggregation (3:l aggregates or still higher aggregates) at high mole fractions of acid are clearly present. Figure 2 shows duE/dT (averaged between 293.15 and 313.15 K) for the same systems. The behavior of uE is parallel. However, it is remarkable that duE/r3T is so strongly negative. This implies that uE has virtually nothing to do with complex concentration (because practically all molecules available are complexed, irrespective of the temperature), but rather with the properties of the complexes. When one takes into account our results on the system acetic acid + trifluoroacetic acid: the sugges-
+
~
~~~~
(6)Kohler, F.; Findenegg, G. H.; Bobik, M. J.Phys. Chem. 1974, 78, 1709.
2528
Kohler et al.
The Journal of Physical Chemistry, Vol. 85, No. 17, 1981
TABLE I: Values o f the Parameters Ai o f Eq 1 and Standard Deviations o [ u E / ( x , x , ) ]o f the Experimental Quantities u E / ( x , x , )for Mixtures Carboxylic Acid ( x , ) t Triethylamine ( x , ) at 293.15 and 313.15 K T/K 293.15 313.15 293.15 313.15 293.15 313.15 293.15 313.15
HCOOH t (C,H,),N (for 0 < x, < 0.4) CF,COOH t (C,H,),N (for 0 < x , < 0.5) C,H,COOH t (C,H,),N (CH,),CCOOH t (C,H,),N
A0 -42.936 -50.173 -57.686 -65.187 -32.093 -34.145 -27.741 -29.544
A, -11.505 -18.278 -39.582 -48.764 29.914 31.664 29.451 31.019
A, 24.676 29.036 4.935 -6.851 -1.297 - 7.400 -14.435 -19.157
-43
18.459 26.471 60.151 69.938 -17.669 - 17.521 -24.369 -23.804
[~E/(x,x,)l
A4 0.0 0.0 6.228 13.887 5.039 10.482 28.177 34.197
0
06
08
0.708 0.438 0.670 0.57 2 0.860 0.637 0.615 0.693
-10
avE
ai
,pCm3 mol K
-50
\,
, 0
02
04
06
08
10
02
0
04
x2
x2
Figure 2. Temperature dependence of the molar excess volume
a vE18T (between 293.15 and 3 13.15 K), for the same systems as in Figure 1. Previous results5 on acetic acid (CH,) shown for comparison.
+ trlethylarnine are
+
+
n-1
[uE/(xlxz)]/(cm3mol-') = C Aj(xl - xdi i=O
(1)
where x2 denotes the mole fraction of amine. Values of the n coefficients Ai and standard deviations u[uE/(xlx2)l are contained in Table I. Because of the complicated behavior of uE/(x1x2), the correlation shows systematic deviations from the experimental values, but it is still acceptable for calculating the density at any composition. In the case of the n-butylamine and the di-n-butylamine systems, a satisfactory representation of the composition dependence of uE/(x1x2) is very difficult to achieve. A good (7) Coleman, R. N.; Prideaux, B. R. J. Chem. SOC.1937,1022.
+
Figure 3. Quantity V ~ / ( X , X of ~ ) the systems propIonic acid dl-nbutylamine (Bu,NH) and propionic acid n-butylamine (BuNH2) at 293.15 K vs. mole fraction of amine (xz). Also shown are the curves for propionic acid triethylamine (Et3N)of Figure 1 and propionic acM diethylamine' (Et,") at 298.15 K for comparison. The dashed portion of the curve of proplonic acid dLn-butylamine is Interpolated through the region of solidification.
+
+
+
tion seems appropriate that the polarization of the complexes increases with temperature. However, in view of the fact that the product K q (discussed below) decreases with increasing temperature, this suggestion can only be made with caution. Finally, Figure 3 gives uE/(x1x2) for propionic acid + n-butylamine and propionic acid di-n-butylamine, with the curves for propionic acid triethylamine and propionic acid + diethylamine' for comparison. As already inferred from viscosity measurements,' the primary and secondary amines show both 3:l aggregation and 1:l complex, though uE/(x1x2) does not exhibit the extreme composition dependence as observed for the viscosities. All of our volumetric results on the triethylamine systems were smoothed by unweighted least-squares polynomials according to
1
+
correlation could be effected for propionic acid tylamine at 313.15 K by
+ n-bu-
3
C A j ( ~-l xJ2'
~1x2
i=O
uE/(cm3 mol-') = 1+
(2)
2
c Bi(X' -
x2)i
i=O
with A. = -37.7883, Al = -15.3641, A2 = -319.8841, A3 = 85.2916, Bo = 1.5120, B1= 1.2723, and B2 = 7.3661. The corresponding standard deviation was u(uE) = 0.0076. But because of the pronounced peak of the uE/(x1x2)curve of the same system at 293.15 K (cf. Figure 3), the same approach would not work well at this temperature. Also, all of our attempts to obtain a reasonable correlation for #(x& of propionic acid di-n-butylamine failed. Conductometric Measurements. Electrolytic conductivities: K , of formic, trifluoroacetic, acetic, and propionic acids + triethylamine are shown in Figure 4. For trifluoroacetic acid + triethylamine, the product KV (viscosities were taken from ref 1)is calculated and presented in Figure 5, whereas for the other systems KV can be obtained from the more extensive measurements of Huyskens et al? It is remarkable that the K q curves look similar for dl systems including formic acid + triethylamine and trifluoroacetic acid triethylamine, whereas the viscosities
+
+
~~
(8) A p r e l i m i i report on these measurements was given by: Kohler, F.; Liebermann, E.; Wilhelm, E. "Proceedings", Symposium on Thermodynamic, Spectral and Structural Features of Interactions between Molecules, Excited Molecules or Ions and Ligands, WOpion/Louvain, Belgium, Sept 16-20, 1974, p 108.
The Journal of Physical Chemistry, Vol. 85, No. 17, 198 1 2527
Mixtures of Carboxylic Acids with Amines
CH, CH,
b
10'~
0.2
0
0.4
0.8
0.6
+
Flgure 7. Product KV for propionic acld n-butylamine (BuNH,), propionic acM dl-n-butylamine (Bu,NH), and propionlc acld tri-nbutylamine (Bu3N)at 293.15 K vs. mole fraction of amlne (x,).
1
+
1
+
x2
Figure 4. Electrolytlc conductivlties, K, of the systems formic (H), trifluoroacetlc (CF,), acetic (CH,), and propionic (CH,CH,) acids triethylamine at 293.15 K vs. mole fraction of triethylamine (x2).
+
I
1
/
/i
00
02
O4
8
:/
YH'
x2
Figure 5. Product of electrolytic conductivity times viscosity. K V , for trlfluoroacetlc acid triethylamine at 293.15 K vs. mole fractlon of triethylamine (x,).
+
3 x __
R-'Cm"
Figure 8. NMR chemical downfield shift, -8, of various protons of trifluoroacetlc acld triethylamine at 313.15 K vs. mole fraction of amine (x,). For details see text. Note the shifts of scale of the ordinate.
+
2 10':
110:
+
+ +
0 Figure 6. Electrolytic conductivities, K, of the systems propionic acid n-butylamlne (BuNH,), propionlc ackl di-n-butylamine (BuaH), and propionic acid trl-n-butylamine(Bu,N) at 293.15 K vs. mole fractlon of amine ( x p ) . The curve for propionic ackl 4- triethylamine (EbN) of Flgure 4 is shown for comparison. The dashed part of the curve of propionic acM di-n-butylamineindicates the reglon of soildiflcation.
+
and the 9values do not. The high KT values of formic acid + triethylamine are caused by the high relative dielectric permittivity of formic acid. We will show later that the 1:l complex of trifluoroacetic acid triethylamine has a zwitterion structure, which is less pronounced in formic acid + triethylamine, and much less with all other systems. Electrolytic conductivities of propionic acid + tri-nbutylamine, propionic acid di-n-butylamine, and propionic acid + n-butylamine are presented in Figure 6, with the curve for propionic acid triethylamine for comparison. The products KV are shown in Figure 7. Note that for mixtures with di-n-butylaminemost of the ionization still takes place around x,,id = 0.75, whereas with n-butylamine the large cluster at %,id = 0.5 carries most ions. NMR Measurements. Figure 8 gives the chemical shift 6 of the acidic proton of trifluoroacetic acid, and the
+
+
+
2528
The Journal of Physical Chemistty, Vol. 85, No. 17, 1981 ,
Kohier et ai.
4 1
I
I
A
04
02
0
1
x2
Flgure 9. As in Figure 8, but for formic acid
+ triethylamine (x2). A
I
1.4 -
I
C$ of Et,N 1.0
'I
\
3.2
C& of E t j N
1 12
-6inppm
i
10
I
0
A
02
04
I
1
x2
Figure 10. As Figure 8, but for acetic acid
+ triethylamine (x2).
methylene and methyl protons of triethylamine, in the mixture trifluoroacetic acid triethylamine. The methylene protons are relatively strongly shifted downfield at high acid mole fractions, and their quartets are doubled, which shows the substantial presence of the acidic proton in the ion +HN(CH2CH& The acidic proton appears in two signals: one is connected with trifluoroacetic acid and can be shifted upfield by adding small amounts of water; the second is broader and is caused by the proton on the ammonium ion in the 1:l complex, which cannot be shifted by adding small amounts of water. The integral value of the two signals adds up to one proton, and the distribution is linear between the pure acid and the equimolar mixture. Comparison between the NMR behavior and the conductivities or the K ~ curves, I respectively, shows that the equimolar mixture, though made up almost exclusively of the 1:l zwitterion, is not the one most conducting. It is rather the 3 1 mixture where the zwitterion polarizes other associates and where, possibly, conductivity is a result of a flipping over of hydrogen bonds and is not due to a migration of whole aggregates. The analogous curves for formic acid triethylamine are shown in Figure 9. The methylene quartets are again split up, but no separate signal was observed for the ammonium ion, in spite of the higher concentration of dissociated ions. In addition, the signal for the methyl protons is shifted somewhat less downfield than in trifluoroacetic acid mixtures, indicating that in this 1:l complex
8
+
+
6
\
4
3.0 2.8 2.6
0
0.2
0.6
0.4
0.8
x2
Figure 12. As Figure 8, but for propionic acid at 296.5 K.
+ n-propylamine ( x p )
the proton is not completely attached to the nitrogen. This casts some doubt on the interpretation of Huyskens et ale: according to which triethylamine in formic acid around Xa&d = 0.75 should be almost exclusively in the form of the ammonium ion. We rather prefer to speculate that in these mixtures conductivity is partly caused by a flipping over of hydrogen bonds, so that Walden's rule is not strictly applicable. Figures 10 and 11 present the NMR signals for acetic acid triethylamine and propionic acid + triethylamine. Though the downfield shift of the methylene protons is
+
J. Phys. Chem. 1981, 85, 2529-2530
almost as strong as in formic acid or trifluoroacetic acid mixtures, a split of the methylene quartets was not observable any more. Further, the downfield shift of the acidic proton becomes somewhat less pronounced than with the stronger acids. The NMR data for the acidic proton in propionic acid + triethylamine have been measured by different groups under different experimental conditions; they are very sensitive with respect to the purity of the acid. The composition dependence of S for the acidic proton as well as for the methylene protons adjacent to nitrogen does not change significantly when triethylamine is replaced by tri-n-propylamine or tri-n-butylamine. However, a drastic change is observed when the tertiary amines are replaced by the primary n-propylamine. This is shown in Figure 12. It is evident that the acid-amine hydrogen bond is less polar. On the other hand, the methylene protons at equimolar composition are more affected than in the systems with tertiary amines. If we compare this behavior with either the conductivity data or the KV curve of propionic acid + butylamine, it seems again hard to believe that ionization is so strong that
2529
ionized aggregates do account for all of the conductivity. Acknowledgment. We are grateful to the Deutsche Forschungsgemeinschaft for support of the density measurements and to the Deutsche Forschungsgemeinschaft and the Office of Cultural Relations of the Hungarian People’s Republic for facilitating the cooperation between Ruhr-Universitat Bochum and Veszpremi Vegyipari Egyetem. Supplementary Material Available: Tables containing the following: (i) values of excess volumes at various mole fractions of amine ( x 2 ) for the systems formic, trifluoroacetic, propionic, and trimethylacetic acids + triethylamine, propionic acid + di-n-butylamine, and propionic acid + n-butylamine, all at 293.15 and 313.15 K; (ii) electrolytic conductivities of formic, acetic, and propionic acids + triethylamine and propionic acid tri-n-butylamine at 293.15 and 313.15 K, of propionic acid di-nbutylamine at 293.15 and 318.15 K, and of trifluoroacetic acid + triethylamine and propionic acid + n-butylamine at 293.15 K (9 pages). Ordering information is available on any current masthead page.
+
+
Vapor-Pressure Lowering of Anhydrous Hydrogen Fluoride by an Involatile Solute Peter McTlgue Department of Physical Chemistty, Universw of Melbourne, Parkvllle, 3052, Australia (Received: March 17, 198 1)
The extent of the lowering of the vapor pressure of anhydrous hydrogen fluoride (AHF) by an involatile solute is “abnormally” large because of the high degree of association in HF vapor. Existing experimental data are used to illustrate this previously unnoted effect.
The activity of a solvent in a solution at some fixed temperature is given by a =f
/f
(1)
where f and f are the fugacities of the solvent vapor in equilibrium with the solution and the pure solvent, respectively. The vapors of most common solvents behave nearly ideally, and (1)may then be replaced to a fair approximation by a = P/P‘
(2)
where p and p’ are the pressures exerted by the solution and the pure solvent, respectively. If, on the other hand, the solvent vapor is highly associated, a satisfactory approximation to the solvent activity would be
main gas-phase equilibria have been identified as 2HF = (HF), 6HF = (HF), Janzen and Bartel13have analyzed all existing data to 1968 and have prepared tables of the mole fraction of monomeric HF at various total HF pressures up to the saturation vapor pressure, from -36 to 24 “C. From (3), then, the €IF activity of any solution will be given by UHF
= PHF/PHF*
(4)
(3) where pmand pm*are the partial pressures of monomer present in the vapor of the solution and the pure solvent, respectively. Many investigations have shown that AHF is virtually unique among ligands in having a vapor that is highly associated even at quite low The
where pm and pm’ are the monomer partial pressures over the solution and pure solvent, respectively. Figure 1shows a plot of UHF calculated in this way, as a function of total HF pressure at 0 “C, using Janzen and Bartell’s tables and the moxt recent4svalue of p’, the saturation vapor pressure of AHF at 0 “C. The data of Figure 1 can be used to calculate the expected total vapor pressures of HF above solutions of both dissociated and undissociated involatile solutes at 0 “C.
(1)Simons, J.; Hildebrand, J. H. J.Am. Chem. SOC.1924,46, 2183. (2)Janzen, J.; Bartell, L. S. J. Chem. Phys. 1969,50, 3611. (3)Janzen, J.; Bartell, L. 5. U.S. Atomic Energy Commission IS-1940, 1968,and references cited therein.
(4) Sheft, I.;Perkins, A. J.; Hyman, H. J. Inorg. Nucl. Chem. 1976,35, 3677. (5) See, e.g. Robinson, R. A,; Stokes, R. H. “Electrolyte Solutions”; Butterworths: London, 1959; 2nd ed, p 34.
a = Pm/Pm’
0022-3654/8 1/2085-2529$01.25/0
0 1981 American Chemical Society
