1882
T. L. Buxton and J. A.
Chlorobenzene shows somewhat different behavior. Plots of 6 M - ,6 US. x for C H ~ C H ~ C O Zand - CH2CO2have similar slopes and those for CN3(CHz)N+ are essentially invariant with z. These results might imply a somewhat deeper penetration by chlorobenzene into a given alkylammonium propionate micelle than by the aliphatic hydrocarbon solvents. The considerably higher dielectric constant, stronger dipole character, and ordered structure of dipolar aprotic DMAC merits separate discussion. Indeed the monomeric and micellar chemical shift data for the alkylammonium propionate surfactants in this solvent possess unique features (Table 111 and Figure 4).Most striking is the behavior of the NH3+ protons; a plot of t 3 ~- 6, us. x has a negative slope! This behavior can be more profitably discussed in terms of comparing changes of ,6 and 6~ as functions of the alkyl chain length (x) in DMAC with those in the other solvents. It is evident that an increase in chain length of the surfactant affects the monomeric chemical shifts considerably more in CDCl3, CH2C12, and CGHsCl than in DMAC while the converse is true for 6M. It is tempting to rationalize these results in terms dipoledipole interactions i n the monomeric species
which are only slightly affected by changes in x. Formation of micelles requires, however, strong desolvation of the monomers and values of values for 6~ will change rapidly as x increases due to the formation of more compact micelles. Changes in the polarity of the solvent are likely to alter the structure of the micelles. A transition from a spherical or ellipsoidal structure to a lamellar or "staggered" one is not improbable with a relatively large increase in solvent polarity. There may be, of course, a solvent polarity range in which micelle formation is not observable. In summary, formation and structures of micelles in nonaqueous solvents are delicately balanced between the interaction of functional groups of the surfactants with each other
Caruso
and with the solvent both in the monomer and in the aggregate. Alterations in the size and geometry of the surfactant functional groups and polarizability as well as such properties of the solvent as polarity, bulk, and geometry are the many factors which may be responsible for the ultimate arrangement of solvent molecules in and around the aggregated surfactant. Efforts in our laboratories are directed toward the systematic elucidation of these factors.
Acknowledgments. Support for this work from the Robert A. Welch Foundation and from the Research Corporation is gratefully acknowledged. E. J. F. is a Research Career Development Awardee of the National Institutes of Health, U. S. Public Health Service. References and Notes (1) Part II: E. J. Fendler, J. H. Fendler, R. T. Medary, and 0 A. El Seoud, J. Phys. Chem., 76, 1432 (1973). (2) J. H. Fendler, E. J. Fendler, R. T. Medary, and 0. A. El Seoud, J. Chem. SOC.,Faraday Trans., 7, 69, 280 (1973). (3) F. M. Fowkes in "Solvent Properties of Surfactant Solutions," K. Shinoda, Ed., Marcel Dekker, New York, N. Y . , 1967, p 65. (4) A. Kitahara in "Cationic Surfactants," E. Jungermann, Ed., Marcel Dekker, New York, N. Y., 1970, p 289, and references cited therein. (5) E. J. Fendler, J. H. Fendler, R. T. Medary, and V. A. Woods, J. Chem. SOC. D. 1497 (1971). (6) J. H. Fendler, J. Chem. Soc.. Chem. Commun.. 269 (1972). (7) J . H. Fendler, E. J . Fendier, R. T. Medary, and V. A. Woods, J. Amer. Chem. SOC.,94, 7288 (1972). (8) C. J. O'Connor, E. J. Fendler, and J. H. Fendler, J. Amer. Chem. SOC.,95, 600 (1973). (9) A. Kitahara, Bull, Chem. SOC. Jap., 28, 234 (1955); A. Kitahara, ibid., 30, 586 (1957). (10) H. InoueandT. Nakagawa, J. Phys. Chem., 70, 1108 (1966). (11) J. C. Eriksson, Acta Chem. Scand., 17, 1478 (1963). (12) N. Muller and R. H. Birkhahn, J. Phys. Chem., 71,957 (1967). (13) N. Muller and R. H. Birkhahn, J. Phys. Chem., 72, 583 (1968). (14) R. Haque, J. Phys. Chem., 72,3056 (1968). (15) R. E. Bailey and G. H. Cady, J. Phys. Chem., 73, 1612 (1969). (16) I . J. Lin and P. Somasundaran, J. Colloid Interface Sci., 37, 731 (1971), and references cited therein. (17) E. M. Kosower, "An Introduction to Physical Organic Chemistry," Wiley, New York, N. Y., 1968. (18) L. M. Jackman and S . Sternhell, "Applications of Nmr Spectroscopy in Organic Chemistry," Pergamon Press, Elmsford, N. Y., 1969; A. D. Buckingham, T. Schaefer, and W. G. Schneider, J. Chem. Phys., 32, 1227 (1960).
Spectroscopic Studies of Solvation in Sulfolane T. L. Buxton and J. A. Caruso* Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 45227
(Received October 30, 1972)
Infrared spectroscopy has been used to study ion solvation in sulfolane, a dipolar aprotic solvent of high (44)dielectric constant. Lithium, sodium, potassium, rubidium, and cesium salts were investigated. The solute absorption bands appear to be independent of anion. The spectral data indicate that solvation takes place in a manner similar to that observed for dimethyl sulfoxide solutions and that sulfolane is a far weaker donor. It is proposed that solvent-separated ion pairs may be formed in sulfolane solutions.
In the past few years there have been a number of spectroscopic studies of solvation in nonaqueous solvents.l-ll The solvents studied have varied from those with low dielectric constants such as tetrahydrofuran,l>9 benzene,ll The Journal of Physical Chemistry. Vol. 77, No. 75. 7973
and pyridines to moderate dielectric constants such as dimethylformamidel0 and the pyrrolidones6 to those of moderately high dielectric constants such as dimethyl sulfoxide.3-5
Spectroscopic Studies of Solvation in Sulfolane
w 0.5
TABLE I Salt
u Concn, M
0.50 0.50 0.50 0.50 0.25 0.25 0.25 0.25 0.25 ~ 0 . 2 5(satd)
5 0.4 g Q3
v(max), c m - ’ 187 f 2 186 f 2 186 f 2 156 f 5 152 f 5 156 i 5 197 f 2 1 9 7 ;+ 2 195 f 2
~ 0 . 2 5(satd) 0.50
0.50
Sulfolane is a dipolar aprotic solvent with a moderately high dielectric constant (44).12 It has been shown to dissolve chlorides and perchlorates to a reasonable extent.13 Conductance studies have indicated that ions are slightly associated and have led to solvation numbers at 2.0, 1.5, 1.4, 1.3, and 0.9 for Na+, K+, Li+, Rb+, Cs+, and NH4+, respectively.14 315 In these laboratories, sulfolane has been investigated as a medium for the study of weak acids16 and for the analysis of drugs.17 This study was undertaken to gain further insight into ionic solvation in sulfolane. Experimental Section Sulfolane (Phillips Petroleum Co.) was purified as previously described.16 Sodium and ammonium iodide (Baker Analyzed Reagents) and sodium perchlorate (G. Frederick Smith) were recrystallized from acetone with ether and vacuum dried. Sodium tetraphenylborate and potassium iodide (Fisher) and potassium perchlorate (G. Frederick Smith) were vacuum dried and used without further purification. Potassium and ammonium tetraphenylborates were synthesized by adding an aqueous solution of the respective iodide to one of sodium tetraphenylborate and washing until iodine free. They were then vacuum dried. Ammonium perchlorate was synthesized by adding perchloric acid to an aqueous solution of the iodide; the precipitate was vacuum dried. Rubidium perchlorate, cesium perchlorate, lithium bromide, and lithium chloride were all the purest grade available from Alpha Ventron and were vacuum dried and used without further purification. Infrared spectra were recorded between 4000 and 200 cm-I using a Beckman Model IR-12 spectrophotometer. Infrared spectra between 400 and 20 cm-I were recorded on an R.I.I.C. (Beckman) FS-720 Fourier spectrometer equipped with FTC 300 Fourier transform electronics. All spectra on the TR-12 were run using a Beckman demountable cell with 0.05-mm Teflon spacers and KBr windows to 400 cm-I; polyethylene windows were used from 400 to 200 cm-1. All spectra on the FS-720 were run with R.I.I.C. FS-03 vacuum-tight cells with polyethylene windows and 0.075-mm Teflon spacers. A Du Pont Model 310 curve analyzer was used to help resolve some of the far-ir spectra. Results a n d Discussion Katon and Feairheller have described the structure of sulfolane and made infrared band assignments.18 Our spectra agree reasonably well with theirs with two excep-
0 g Q2
0.1
0 01 0.2 03 0.40.5 0.6 MOLARITY ( N a I ) Relationship of absorbance to concentration for solute-solvent band using Nal. Figure 1.
tions. They report no bands below 242 cm-1 although the spectra were run to 80 cm-1. Below 100 cm-1 we find a strong band that is probably due to dipole-dipole interactions as noted for other polar solvents.1g A second extremely weak band was seen with a center about 140 cm-1 and was partly overlapped by the dipole-dipole band. This may be the band due to C-S-C ring puckering which Katon and Feairheller predicted but did not observe.18 The spectral tabulation for salt solutions is given in Table I. It was originally anticipated that all solutions would be run at 0.5 M ; however, solubility difficulties precluded this for some salts. The theoretical reproducibility of the band maxima for the FS-720 is i 0 . 5 cm-1, but because the bands are broad and often overlap the dipole-dipole band, the best that could be achieved was f 2 cm-1. The potassium peak is almost totally overlapped by the dipole-dipole band and is therefore much more difficult to locate accurately. No bands are seen for rubidium or cesium salts, probably because they fall at lower energy and are hidden by the dipole-dipole band. Neither is a band seen for lithium salts, This band maximum would probably come at about 400 cm-1 as in other solvents.5-8Jo Unfortunately, sulfolane has strong absorption bands between 440 and 310 cm-1. Attempts to find an absorption band for the lithium solutions up to 1.0 A4 were made. No bands were found and no change could be found for any of the other bands in the sulfolane spectra. The solute absorption bands observed have a linear relationship between absorbance and concentration. Figure 1 is an example with NaI as solute. The bands seem to be independent of the anion, which would be expected in a polar solvent like ~ u l f o l a n eThere .~ was no shift in the peak location a t different salt concentrations. Sulfolane and dimethyl sulfoxide have been compared previously as solvents.20,21 In spite of similar dielectric constants, 44.0 for sulfolanef2 and 46.4 for dimethyl sulfoxide6 and the S=O group through which the molecule probably interacts with solute, sulfolane has consistently shown weaker interactions with ionic species.20.21 The Gutmann donor numbers also reflect this order, 14.8 for sulfolane22 and 29.8 for Sulfolane’s far infrared solvent-cation bands were at lower energy and weaker than those in DMS0.3.5 This indicates that solvation takes place in a similar manner in both solvents but that sulfolane is a far weaker donor. Previous studies have shown evidence for ion pair formation in sulfolane.14J5~20 The anion independence of the far-ir band shows that it does not enter into the inner solvation shell and therefore The Journal of Physical Chemistry. Vol. 77. No. 15. 1973
Barbara J. Barker and Joseph A . Caruso probably does not form contact ion pairs. This does not rule out the formation of solvent-separated ion pairs as has been postulated for other solvents.24 It seems probable that in sulfolane this type of pairing occurs. References and Notes (1) W. F. Edgell, J. Lyford, I V , R . Wright, W. Risen, Jr., and A. Watts, J. Amer. Chem. SOC.,92, 2240 (1970). (2) E. G. Hohn, J. A. Olander, and M. C. Day, J. Phys. Chem., 73, 3880 (1969). (3) 8. W. Maxey and A. I. Popov, J. Amer. Chem. Soc., 89, 2230 (1967). (4) B. W. Maxey and A. I . Popov, J. Amer. Chem. SOC., 90, 4470 (1968). (5) 5.W. Maxey and A. I. Popov, J. Amer. Chem. Soc., 91, 20 (1969). (6) J. L. Wuepper and A. I. Popov, J Amer. Chem. SOC., 91, 4352 (1969). (7) J. L. Wuepper and A. I . Popov, J , Amer. Chem. SOC., 92, 1493 (1970). (8) W. M. McKinney and A. I . Popov, J. Phys. Chem., 74, 535 (1970)
(9) W. F. Edgell, A. J. Watts, John Lyford, I V , and W. M. Risen, Jr., J. Amer. Chem. SOC.,88, 1815 (1966). (10) c. Lassigneand P. Baine, J. Phys. Chem., 75, 3188 (1971). (11) J. C. Evansand R. Y . Lo, J. Phys. Chem., 69, 3223 (1965). (12) E. M. Arnett and C. F. Douty, J. Amer. Chem. SOC.,86, 409 (1964), (13) J. A. Starkovich and M. Janghorbani, J. inorg. Nuci. Chem., 34, 789 (1972). (14) M. Della Monica, U. Lamanna, and L. Senatore, J. Phys. Chem., 72, 2124 (1968). (15) M. Della Monicaand U. Lamanna, J. Phys. Chem., 72 4329 (1968). (16) P. M. P. Eller and J. Caruso, Anai. Lett. 4, 13 (1971). (17) T. Buxton and J. Caruso, Taianta, 20, 254 (1973). (18) J. E. Katon and W. R. Feairheller, Jr., Spectrochim. Acta, 21 199 (1965). (19) R. J. Jakobsen and J. W. Brasch, J. Amer. Chem. Soc., 86, 3571 (1964). (20) R. Garnsy and J. E. Prue, Trans. FaradaySoc., 64, 1206 (1968). (21) C. Agami and P. Gondouin, Buil. SOC.Chim. Fr., 3930 (1969). (22) V. Gutmann and A. Scherhaufer, Monatsh. Chem., 101, 3930 (1970); 99, 335 (1968). (23) V. Gutmann and U. Mayer, Monatsh. Chem., 101, 3930 (1970). (24) R. H. Ehrlich and A. I. Popov, J. Amer. Chem. Soc., 93, 5620 (1971).
Transport Property Investigation of Ammonium and Tetraalkylammonium Salts in 1,1,3,3-Tetramethylureal Barbara J . Barker and Joseph A. Caruso* Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 45227 (Received October 30, 7972)
The behavior of ammonium and a series of tetraalkylammonium salts in a relatively new, aprotic solvent of moderate dielectric constant, tetramethylurea (TMU), was studied by conductance and viscosity techniques. Conductance data for all salts were evaluated both by the Fuoss-Shedlovsky method of analysis and by the Fuoss-Onsager equations for unassociated and associated electrolytes. The effect of tetraalkylammonium salts on the viscosity of TMU was determined by experimentally evaluating viscosity B coefficients of the Jones-Dole equation; the effect of solution viscosity on the electrolyte size parameter C ~ Jof the Fuoss-Onsager theory was determined by including viscosity corrections in the conductance data evaluation. Tetraalkylammonium perchlorates and bromides were found t o be associated electrolytes in TMU. Ionic limiting equivalent conductances in TMU were obtained indirectly by using triisoamylbutylammonium tetraphenylborate as a reference electrolyte. The order of decreasing relative cationic limiting equivalent conductances which was observed in tetramethylurea also has been noted in several other nonaqueous solvents.
Introduction Frequent use of nonaqueous solvents as media for organic and inorganic reactions has led to increasing interest in the transport properties of electrolytes in solution. Electrolyte behavior in solution is revealed by fundamental viscosity and conductance investigations which elucidate directly from experimental measurement the nature of ion-ion and ion-solvent interactions. Recently there have been a number of transport property investigations in polar nonaqueous solvents such as propylene carbonate (PC),2 dimethyl sulfoxide (DMS0),3-9 sulfolane,6 acetonitrile (ACN),7-Qacetone,lOJl methyl ethyl ketone,l2 and aliphatic a l c ~ h o l s . l ~Tetramethylurea -~~ (TMU), a relatively new aprotic solvent of moderate dielectric constant (23.45) and wide liquid range (from -1 to 176.5"), proved to be an interesting nonaqueous medium for a recent conductance investigation16of alkali metal salts. The Journal of Physical Chemistry, Voi. 77. No. 75. 7973
Quaternary ammonium salts often are studied as theoretical models in solution chemistry since they form systematic electrolyte series containing cations of relatively large crystallographic size and low charge density. Since tetramethylurea promises to be a solvent of considerable interest and since no transport property investigations of tetraalkylammonium salts in 'I'MU have been reported, the present investigation was undertaken. Experimental Section The purification of tetramethylurea has been described previously.16 Ammonium tetraphenylborate was prepared by mixing equimolar aqueous solutions of ammonium bromide (Fisher Certified Reagent) and sodium tetraphenylborate (Baker Analyzed Reagent). The precipitate was recrystallized from an acetone-ether mixture. Tetramethylammonium perchlorate was prepared by adding an ap-
