Oct., 1958
PHYSICAI PROPERTIES O F THE
SYSTEM ETH24NOL-2,2,2-TRIFLUOROETHANOL
1311
PHYSICAL PROPERTIES AND HYDROGEN BONDING IN THE SYSTEM ETHANOL-2,a,2-TRIFLUOROETHANOL1 BY L. M. MUKHERJEE AND ERNEST GRUNWALD Chemistry Department, Florida State University, Tallahassee, Fla. Received February 8 , 1968
A number of physical properties of the system ethanol-2,2,2-trifluoroethanol have been measured: the dielectric constant, density, viscosity and their temperature coefficients. The dielectric constant varies by only 6% throughout the entire system. The density and viscosity indicate that the system is very non-ideal. The boiling point-composition diagram was determined and shows a maximum boiling point (81.75') a t 57.65 wt. yo trifluoroethanol. The ~ K ofA trifluoroethanol in water at 25' was redetermined and found to be 12.5 It 0.1, as compared to the value 16.0 for ethanol. Hydrogen bonding was studied by infrared spectrometry in the range 3800-3200 cm.-l for dilute solutions of the alcohols and their mixtures in carbon tetrachloride. The infrared spectra of trifluoroethanol at various concentrations resemble those of ethanol and point to the formation of dimers as well as higher complexes. The association of ethanol with trifluoroethanol is considerable, the association constant for the formation of an 1: 1 complex being estimated as 10 ( M - 1 ) . The viscosity data for the system ethanol-trifluoroethanol are consistent with a model in which ethanol is the terminal acce tor and trifluoroethanol the terminal donor of chain-like mixed complexes. The partial molal volumes are strongly affectez by the complexing.
Ethanol (EtOH) and 2,2,2-trifluoroethanol (TFE) resemble each other in their molecular structures, boiling points and dielectric constants but differ considerably in acidity, basicity and solvating ability. Since the two liquids are miscible in all proportions, the system EtOH-TFE is of special interest for the study of solvation and other properties of dissolved electrolytes, particularly since the long-range interionic forces and Born charging energies are nearly the same a t every composition. The purpose of this investigation is to measure some physico-chemical properties that are useful in the study of electrolytes, and some additional properties in order to illuminate the hydrogen bonding among the components.
Experimental Materials.-TFE, obtained from the Pennsalt Chemicals Corp., was dried over anhydrous magnesium sulfate and a little potassium carbonate, and then distilled through a 15-plate column. The middle fraction, b.p. 73.75" (760 nim.), was used in this work. A gas chromatogram indicated 5 bands: the major band was due to TFE; the other bands were due to dissolved air (area 0.64), water (area 4.9), ethanol (area 2.0) and an unidentified substance (area 2.9). Assuming that the solubility of air in T F E is of the same order as that in EtOH (0.02 wt. yo),the total impurities amount to about 0.3 wt. yo. Commercial absolute alcohol was dried by the method of Mukherjee.* Karl Fischer titration and density determination indicated a water content of about 0.1% for the EtOH used in this work. Dielectric Constants.-Dielectric constants were measured at 20 and 25' by the resonance method38 using a Boonton &-meter, a General Radio type 722-N calibrated variable condenser, and a conventional cell of three concentric platinum cylinders (purchased from J. C. Balsbough, Cambridge, Mass.). The cell was calibrated with air and ethanol,* and dielectric constants D were computed from the equation C = a bD, where C is the observed capacitance, and a and b are constants for the cell. Densities.-Measurements were made with an accuracy
+
(1) Work supported by the Office of Naval Research. Reproduction in whole or in part is permitted for any purpose of the United States Government. (2) L. M. Mukherjec, Sci. and CuEture ( I n d i a ) , 19, 314 (1953). (3) "Physical Methods of Organic Chemistry," A. Weissberger. editor, Interscience Publishers, New York, N. Y., 1949: (a) C. P. Smyth. "Determination of Dipole Moments," Chapter 24; (b) N. Bauer, "Density," p. 264; (c) T. E. McGoury and H. Mark, "Viscometry," p. 334. (4) G. Akcrlof, J. Am. Chem. Soc., 54, 4125 (1932); see also data b y 3 . L. Hall and H. Phillips, communicated t o E. F. Sieckinan and E. Grunwald, ibid.. 7 6 , 3855 (1954).
of 0.0570, using a bicapillary pycnomoeterab which was calibrated with water a t both 25 and 35 Viscosities.-An Ostwald viscometer was used. It was calibrated with water at both 25 and 35", and the liquids under measurement were protected from atmospheric moisture with a guard tube filled with Drierite. (The guard tube had no effect on the time of flow.) The coefficients of viscosity were computed without any correction for capillary end-effects.80 The standard deviation of duplicate measurements was 0.4%. Boiling Point-Composition Diagram.-Liquid-vapor equilibrium was established in a Cottrell apparatus by standard techniques.6 Equilibrium was attained in about 10 min. Boiling points were measured on calibrated thermometers and were corrected for exposed stem and to a pressure of 1 atm. In making the small corrections for pressure difference, d P / d T was assumed to be tJhe same as for ethanol.6 Compositions of the vapor condensates and of the liquid phases were determined by density measurement a t 25", using 0.4-ml. samples and a calibrated Agla Microsyringe. ~ K Determination.-Due A to discrepancies in previous value^,^ ~ K for A T F E was redetermined in water at 25'. A differential potentiometric technique was used,8 employing an Ag-AgC1 and a high pH glass electrode, and a Beckman model GS pH meter. Ionic strengths were of the order of 0.01 molar, and all ionic activity coefficients were assumed equal to those for HCl a t the same ionic strength. In calculating ~ K Acorrection , was made for the hydrolysis of sodium trifluoroethylate. Infrared Measurements.-The infrared spectra of dilute solutions in carbon tetrachloride of EtOH, T F E and of their mixtures were measured with a Perkin-Elmer model 21 spectrometer with sodium chloride optics a t room temperature (25'). Matched sodium chloride cells were used, one containing the solution and the other the pure solvent, and the "base line" was checked frequently by having both cells contain carbon tetrachloride. The measurements of association constants were based on optical densities in the region of the fundamental OH-stretching vibration, 37003300 cm.-l.
.
Results System Ethanol-TrifluoroethanoL-Dielectric constants are listed in Table I. The dielectric constant is quite constant; the total variation for the entire system is only 6%. Densities p and coefficients of viscosity 9 are listed in Table I1 and (5) J. W. Rogers, J. W. Knight and A. R. Choppin. J . Chem. Ed., 24, 491 (1947). (0) J. A. Barker, I. Brown and F . Smith, Diec. Faraday Soc., 16, 142 (1953). (7) (a) A. L. Henne and R. L. Pelley, J . Am. Chem. Soc., T4, 1420 (1952); (b) E. T. McBee, W. F. Mareluff and 0. R. Pierce, ibid., 7 4 , 444 (1952); (0) C. W. Roberts, E. T. McBee and C. E. Hathaway, J . Or& Chem., 21, 1369 (1956). (8) A. L. Bacarella, E. Grunwald, H. P.Marshall and E. L. Purlee, %bid.,20, 747 (1955).
L. M. MUKHERJEE AND ERNEST GRUNWALD
1312
Vol. 62
behavior are shown very clearly by the partial molal volumes of the components (see Fig. 1). Although TFE is more viscous than EtOH, addition of TFE to EtOH lowers the viscosity until a minimum value is reached a t about 40 wt. % TFE. The temperature coefficients of density and viscosity appear to be monotonic functions without significant abnormalities. The deviation from ideal behavior is also shown clearly by the boiling point-composition diagram, Fig. 2. p LI P ~ K Values A in Water.-pKA for ethanol in water I has been estimated as 16.0,9 p& for trifluoroethanol in water is appreciably smaller; however, these several values have been reported previously: 11.40,7a12.30,7b12.43.7c We have redeterIb mined ~ K for A trifluoroethanol and have obtained the value 12.5 rt 0.1 at 25", in adequate agreement with the most recent of the previous values. Association of Ethanol and Trifluoroethanol in trifluoroethanol is a --1216 0 0.5 1.0 Carbon Tetrachloride.-Since stronger acid than ethanol, it is probably also a better donor in hydrogen bonding. In order t o Mole fraction of TFE. obtain more direct evidence we have examined the Fig. 1.-Plots of - Tro for the s stern ethanol-trifluoro- infrared absorption in the region of the fundamental ethanol at 25 : 0, ethanol; trifluoroethanol. OH-stretching vibration. This region is particuindicate that the system is far from ideal. Using larly useful as a specific indicator for hydrogen the density data, deviations from ideal solution bonding since the OH-stretching frequency is lowered slightly in the H-bonded complex.1° The TABLE I self-association of ethanol has been studied in this DIELECTRIC CONSTANTS FOR THE SYSTEM way." There is evidence for the formation of ETHANOL-TRIFLUOROETHANOL" dimers and higher complexes; K2 for dimer formaWt. % tion is 0.64. TFE Daao D 2ao - d log D / d T The infrared spectrum of TFE bears a strong 0.00 (25.00) 24.32 0.0024 f 0.0003 qualitative resemblance to that of EtOH. At 9.64 25.0 24.1 ... concentrations below 0.02 molar, the only signifi22.66 24.8 24.3 ... cant absorption in the OH-stretching region is 25.32 25.1 ... ... that of the monomer a t 3620 f 10 em.-'. At 0.05 50.14 25.2 25.1 ... molar, a new band at 3480 crn.-', analogous to the 57.65 26.06 25.63 0.0014f .0003 lldimer" band of ethanol, is already quite notice73.49 26.0 25.7 ... able. At still higher concentrations there appear 73.68 ... 25.7 ... additional bands a t still lower frequencies, pre100.00 26.53 26.14 0.0013 f ,0003 sumably due to higher complexes.12 Sample spec=Dielectric constants are accurate to 1% if listed to 1 decimal place, and are accurate to 0.2% if listed to 2 decimal tra of ethanol and trifluoroethanol are compared in Fig. 3. The strong analogy between the two places. alcohols suggests that self-association in dilute TABLE I1 solution is of comparable importance. However, 1)ENSITIES AND COEFFICIENTS OF VISCOSITY FOR THE we could not decide from the spectra whether the SYSTEM ETHANOL-TRIFLUOROETHANOLself-association of TFE involves hydrogen bonds only of the type 0-H. . .O, or also of the type Wt. % - (k)" tlr30 E"d
-.h'
f 6,
TFE
0.00 1.945 4.120 4.240 6.210 7.98 8.33 12.11 21.53 28.07 43.07 48.90 73.73 90.10 100.00
" (6p/6T) p o l*
Pao
PT
(cp.)
(kcal.)
0.7856 0.00085 1.102 3.32 .7941 .00089 ... ,8026 .00103 (1) . . . *.. .8032 .00091 1,104 3.44 .8100 .00090 ... ... .8166 .00092 1.102 3.39 .8190 .00092 ... ... .8332 .00101 1.098 3.41 .8703 .OOllO 1.097 3.40 .8998 .00116 1.089 3.45 .9633 ,00114 1.082 3.50 .9900 .00111 1.105 3.50 1.1390 .00140 1.215 3.85 1.2761 .00161 1.427 3.84 1.3816 .00164 1.768 4.55 (pas0 - p26")/10. E,,is0 =: 42.025 [log qzS0/
...
0-H.. .F.
In mixtures of EtOH and TFE, there is strong additional absorption with a maximum at 3400 cm.-1, as shown in Fig. 3c. This is ascribed to the formation of mixed hydrogen bonded species, the most important of which is expected to have the structureF &CH20H.. .0HCH2CHs. At the same time the optical density of the llmonomer" band (9) E. F. Caldin and G. Long, J . Chem. SOC.,3737 (1954). I n order to convert their value to a conventional "molar" basis, the Henry's law constant for ethanol was taken from H. Goller and E. Wicke. Anpew. Cham., B19, 117 (1947). (10) M.M.Davies, Ann, Rep. Prop. Cham., 43,5 (1946); L. Kellner, Rep. Prop. Phgs., 16, 1 (1952). (11) W. C. Coburn and E. Grunwald. J . Am. Chem. Soc., 80, 1318 (1958).and previous references cited in this paper. (12) R. N. Haszeldine, J . Chem. Boc., 1757 (1953), has reported infrared speotra for TFE in carbon tetrachloride but appears to have missed the "dimer" band at 3480 cm.-1 See, however, Fig. 3b.
Oct., 1058
P H Y S I C A L P R O P E R T I E S O F THE S Y S T E M
ETHANOL-2,2,2-TRIFLUOROETHANOL
1313
82 is considerably less than the sum of the optical densities of the separate components at the same concentration. It is not possible to calculate the association constant from the optical density of 80 the “monomer” band without making some arbitrary assumptions. Assuming an 1:1 complex, the minimum value of K is obtained by supposing that the absorption of the complex is negligible a t the maximum of the “monomer” band. K values 9 78 computed in this way are reasonably constant, as 4 shown in Table 111. (The self-association of the F4 two alcohols was neglected in these calculations.) 76 Other methods of treating these data lead to K values as large as 15. K may be obtained in a more straightforward way, but with lower precision, from the optical 74 density of the complex in the neighborhood of I 3400 em.-’. Using a method of calculation described previou~ly,’~the lower confidence limit 0 0.5 1.0 for K was found t o be 10. The only assumption in Wt. fraction of TFE. this calculation is that of an 1:l complex. Thus Fig. Z.--Boiling qoint-composition diagram of the there is no doubt that the equilibrium constant for system ethanol-trifluoroethanol a t 1 atm. : The maximum the formation of a mixed complex is an order of boiling point, 81.75”, occurs at 57.65 wt. TFE. magnitude larger than that for the dimerization of EtOH, and very probably also of TFE.
TABLE 111 INFRARED MEASUREMENT OF ASSOCIAT~ON CONSTANT FOR TFEeEtOH IN CARBON TETRACHLORIDE AT 25” N
-Optical
-Molar EtOH
... 0.1012 ,1012 .I012 .0679 .0340 .0170
conon.-
TFE 0.0904 ,0908 ,0454 .0227 .0800
,0900 .0900
Mixture
...
0.243 ,181 ,147 ,231 ,227 ,227
density‘Separated Molar compoconcn. of nentsb complex 0.214& 0.01 0.338 0.0268 ,236 ,0154 ,0104 .184 .318 .022r ,268 ,0108 ,245 ,0047
...
Mean
K (Y-1)
...
5.6 6.0 9.3 7.3 5.7 4.4 6.4& 1.6
At maximum of “monomer” OH-stretching band at 3020 f 10 cm.-l. Cell length 0.0403 cm. b At the same concentration. a
Discussion I n interpreting the physical properties of the system ethanol-trifluoroethanol, one must recognize the self-association of the two alcohols as well as the strong tendency toward mixed association. Regarding self-association, it is well known that ethanol is a hydrogen-bonded 1 i q ~ i d . l ~The evidence includes not only infrared spectral data, but other physical properties as well, such as the unusually high boiling point and entropy of vaporization (26.8e.u.), well in excess of that expected from Tiouton’s rule. For liquid trifluoroethanol, hydrogen bonding is indicated by infrared spectral data as well as by a high entropy of vaporization (28.0 e.u. at the boiling point15), and the boiling point (73.75’) is higher by 44” than that of the isomeric methoxymethforane, CFaOCH3. The acidbase properties of the two alcohols and their hydrogen-bonding in carbon tetrachloride solutions (13) E. Grunwald and W. C. Coburn, J . Am. Chem. Soc., 80, 1322 (1958). (14) L. Pauling, “Nature of the Chemical Bond,” Cornell University Prees, Ithace, N. Y., 1944.pp. 305,32G. (15) Pennsalt Chemicals Corp., “Trifluoroethanol,” Booklet DC1254,Philadelphia, Ponna., 1956.
3500 3300 Cm.-1 Fig. 3.-Sample infrared absorption spectra a t 25” in the region of the OH-stretchin vibration: (a) 0.1012 molar EtOH; (b) 0.0908 mole %FE; (c) 0.1012 molar EtOH and 0.0908 molar T F E . dashed curve, (a) b) The ordinate is the optical den.& per cm. cell length. $he origin has been displaced vertically in the figures, as indicated. 3700
+
lead one to expect that TFE is a far better “donor” than EtOH, and that EtOH is a far better “acceptor” than TFE. Hence when TFE is added to EtOH, the mixed complexes are likely to be Hbonded chains of the type CFB I CK
CH,
in which TFE is the terminal “donor.” Similarly, when EtOH is added to TFE, one would expect the formation of H-bonded chains in which EtOH is the terminal “acceptor.” In both cases, the
1314
Y. MARCUS
pronounced formation of mixed complexes will affect the self-association equilibria in such a way that the average degree of self-association is lowered. On this basis one can understand why addition of either alcohol to the other results in a lowering of the viscosity. When TFE is added to EtOH, the viscosity decreases only slightly because the greater molecular weight of TFE largely compensates for the decrease in average association number, and the mass of the average kinetic unit appears to remain almost constant. When EtOH is added to TFE, these effects work in the same direction and the viscosity decreases markedly.
Vol. 62
The very complex behavior of the partial molal volumes (Fig. 1) must remain uninterpreted at this time. It is worth noting, however, that the partial molal volume of very dilute TFE in EtOH is markedly smaller than the molar volume of the liquid TFE, while the partial molal volume of very dilute EtOH in TFE is larger than the molar volume of liquid EtOH. It is also worth noting that the minimum in the viscosity occurs a t about the same mole-fraction of TFE, 0.24, as the extrema in the partial molal volumes. On the other hand, the maximum boiling point occurs a t 0.386 mole fraction of TFE.
THE OXIDATION-REDUCTION COUPLES U(1V)-U(V1) AND Fe(I1)-Fe(II1) I N PHOSPHORIC ACID SOLUTIONS BY Y. MARC US^ Contributionfrom the Israel Abmic Energy Commission Laboratories, Rechouot, Israel Received February 68,1868
The otentials of the U(1V)-U(V1) and the Fe(I1)-Fe(II1) couples in 1.2 to 7.5 M phosphoric acid were measured a t 25' anfformal potentials E' were calculated. Plots of E'u and E ' F vs. ~ log U H ~ P O ~the , phosphoric acid activity, were linear, with slopes of (+3.2 =k 0.2) X 0.60/2 mv. and (-1.3 f 0.1) X 60 mv., res ectively. For 1 M phosphoric acid the values are E'u = -475 & 10 mv. and E ' F ~= -555 f 10 mv. us. the normal h Xrogen electrode. Solutions of U(IV), in sharp contrast to Fe(II), in phosphoric acid are very stable to air oxidation. $he behavior of iron and uranium ions, singly or mixed, in phosphoric acid solutions was found to be as might be predicted from the single formal potentials. At hosphoric acid concentrations above 2.4 M Fe(I1) reduces U(VI), and at concentrations below 1.2 M Fe(II1) oxidizes U ( 1 3 .
Recently, Baes2 has published -data on the oxidation-reduction potential of the Fe(I1)-Fe(111) couple in phosphoric acid, and on the oxidation-reduction equilibrium between this and the U(1V)-U(V1) couple. His solutions, however, contained 0.36 M sulfuric acid in addition to phosphoric acid. Previously, Bock and Herrmann3 had measured the potential of the iron couple in phosphoric acid both with and without sulfuric acid, and obtained results which were not in agreement with those of Baes, nor with the earlier results of Carter and Clews.4 Clarification of the issue is desirable; in particular it would be desirable to obtain values for the uranium couple, measured in the absence of sulfuric acid.5 Measurements were made on the uranium and iron couples separately and in mixed solutions, the only anion present being that of phosphoric acid. A few equilibrium measurements in 6 n/r phosphoric acid were also made, for comparison with values calculated from the potentials. Consider the oxidation-reduction process of the iron couple in phosphoric acid where the Fe(I1) and the Fe(II1) species are complexed by phosphoric acid. The position of the equilibrium will depend on the activity of phosphoric acid and hydrogen ions. (1) Taken from the thesis submitted by the author to the Hebrew University, Jerusalem, 1955,for the degree of Ph.D. (2) C. F. Baes, Jr., J . Phys. Chem., 60,805 (1956). (3) R. Bock and M. Herrmann, 2. anorg. allgem. Chem.. 273, 1 (1953). (4) S. M.Carter and F. H. Clews, J . Chem. Soc., 126, 1880 (1924). (5) The work described here waa carried o u t in Israel in 19541955.before the results of Brtes, U. S. Atomic Energy Comm. AECD3594 (19531,cf., Nucl. Sei. Abstr., 8, 190 (1954),became known.
Fe(I1)
+ zH3POrI _ Fe(II1) + e- + gH+
(1)
For a constant phosphoric acid concentration, provided the iron concentration is small compared to the phosphoric acid concentration also, the hydrogen ion activity will be constant, and the potential (in mv. at 25") will be given by
where [ ] are stoichiometric concentrations and y are stoichiometric activity coefficients. Evaluating the expression E - 60 log [Fe(III)I/ [Fe(II) ] from measurements a t constant phosphoric acid Concentration will yield values of E', the formal potential of the couple, defined as its potential vs. the standard hydrogen electrode when the ratio of the stoichiometric concentrations of the oxidized and reduced forms of the substance equals unity E' = E
- 60 log [Fe(III)]/[Fe(II)]
(3)
For the analogous uranium oxidation-reduction reaction, the formal potential at constant phosphoric acid concentrations will be given by E'
=
E
- (60/2) log [U(VI)l/[U(IV)I
(4)
The same situation exists for the oxidationreduction reaction between the uranium and the iron couples U(V1) + 2Fe(II) + xH3P04I_ U(1V)
+ 2Fe(III) + zH+
(5)
for which an apparent concentration quotient may be calculated for every phosphoric acid concentration from the formal potentials
