1296
J. Phys. Chem. 1985, 89, 1296-1304 theories of colloidal interactions in concentrated dispersions.
absorption of light in the latex, the diffracted light has a peak wavelength that corresponds to the spacing near the electrode. This is because the particlelight interaction is so strong that the amplitude of the incident light diminishes rapidly in the latex and particles far from the viewing side do not contribute to the diffraction intensity. In principle, experiments of the kind described in this chapter could yield important information to test various
Acknowledgment. The authors thank Professor S. G. Mason for valuable encollragement. They ah0 are indebted to Professor s. Hachisu for kindly supplying latex samples and for valuable discussions. Registry No. Polystyrene (homopolymer), 9003-53-6.
A Lewis Basicity Scale for Nonprotogenlc Solvents: Enthalpies of Complex Formation with Boron Trlfluorlde In Dlchloromethane Pierre-Charles Maria and Jean-Franqois Gal* Laboratoire de Chimie Physique Organique, Universite de Nice, Parc Valrose-06034 Nice Cedex. France (Received: July 30, 1984; In Final Form: November 2. 1984)
A solvent Lewis basicity scale was established for 75 nonprotogenic solvents by measuring calorimetrically their enthalpies of complexation with boron trifluoride ( A P B F , ) in dichloromethane. Absence of side reactions was verified by calorimetry, spectroscopy, and by checking the stoichiometry of the adducts. Some enthalpies were also measured in nitrobenzene, showing that dichloromethanedoes not induee nonregular effects. Drawbacks of the Gutmann's DN scale are emphasized. Relationships between various Lewis and hydrogen bond basicity scales and - A W B F 1 are examined. A plot of Kamlet-Taft's fl vs. -AHoBF, shows a typical family dependence. A significant multilinear correlation of -AHoBFlagainst complexation enthalpies toward p-fluorophenol and iodine gives evidence that BF3, though stronger, exhibits an electrostatic-covalent acceptor character median between those of the two acids chosen as references. Attention is drawn to the BF, complexation sensitivity to steric hindrance. The - A P B F 3 scale appears as a useful tool for the rationalization of the Gibbs energies of transfer of alkali metal cations which depend mainly on the solvent Lewis basicity. In the correlation analysis of solvent effects the authors suggest the use of basicity parameters representative of the solute-solvent interaction under scrutiny.
Introduction Basicity is an essential solvent property often used to account for the influence of the solvent on chemical phenomena.' Basic solvents are usually classified as HBA (hydrogen bond acceptors)2 or EPD (electron pair donor^).^ Gutmann4 uses the term "donicity" as a measure of the ability to donate an electron pair and he has proposed the so-called donor number (DN) to express, in at least a semiquantitative manner, the donor strength of a solvent. Despite vigorous criticisms of either the concepts or the experimental values: DN is one of the most widely used empirical parameter of solvent basicity, probably by reason of the wide publicity made by the author.' In his excellent review: Jensen brings out the prominent features of the D N concept. We discuss below some failings of the DN scale and we propose another empirical parameter of solvent Lewis basicity defined as the enthalpy change for the reaction between (1) Benoit, R. L.; Louis, C. "The Chemistry of Nonaqueous Solvents", Lagowski, J. J., Ed.; Academic Press: New York, 1978; Vol. 5 , pp 63-1 19. (2) Kamlet, M.J.; Taft, R.W. J . Am. Chem. SOC.1976, 98, 377-383. (3) (a) Reichardt, C. 'Solvent Effects in Organic Chemistry"; Verlag Chemic: Weinheim, 1979. (b) Reichardt, C. Angew. Chem., I n f . Ed. Engl. 1979, 18, 98-1 10. (4) (a) Gutmann, V. 'Coordination Chemistry in Non Aqueous Solutions"; Springer Verlag: Wien, 1968. (b) Gutmann, V. 'The Donor-Acceptor Approach to Molecular Interactions"; Plenum Press: New York, 1978. ( 5 ) (a) Drago, R.S.Coord. Chem. Rev. 1980,33, 251-277. (b) Drago, R. S . Pure Appl. Chem. 1980, 52, 2261-2274. (6) (a) Taft, R. W.; Pienta, N. J.; Kamlet, M. J.; Arnett, E. M. J. Org. Chem. 1981,46,661-667. (b) Olofsson, G.; Olofsson, I. J. Am. Chem. SOC. 1973.95.7231-7233. (c) Lim, Y. Y.; Drago, R.S. Inorg. Chem. 1972,11, 202-204. (7) (a) Mayer, U.; Gutmann, V. Struct. Bonding (Berlin) 1972, 12, 113-140. (b) Gutmann, V.; Schmid, R. Coord. Chem. Rev. 1974, 12, 263-293. (c) Gutmann, V . Coord. Chem. Rev. 1976, 18, 225-255. (d) Gutmann, V. Electrochim. Acta 1976, 21, 661-670. (e) Gutmann, V. Chemfech 1977, 7, 255-263. (0 Gutmann, V.; Resch, G. Comments Inorg. Chem. 1982, 1 , 265-218. (8) Jensen, W. B. Chem. Rev. 1978, 78, 1-22.
0022-3654/85/2089-1296$01 S O / O
gaseous boron trifluoride and the basic organic molecules, including those considered as "solvents", diluted in dichloromethane sol~tion:~ BF,(g)
+ :B(soln)
-
B:BF,(soln)
(1)
Experimental Section Chemicals. Boron trifluoride (Matheson, purity > 99.5%) is purified by slow passage through a trap cooled to about 160 K and then by a freeze-pumpthaw cycle. Every BF, transfer is conducted in a high vacuum apparatus, without contact with the atmosphere, and the gas is stored in a 1-L glass bulb linked to the vacuum line and closed by mean of a mercury valve. Dichloromethane (Merck Uvasol) is stored over a 4-A molecular sieve (Merck) in dark bottles under a blanket of ultrapure argon and used without further purification.1° The water content was measured by Karl Fisher titration and was found to be less than 20 ppm (by weight). Nitrobenzene (Fluka puriss, 99.5% min) was treated in the same way as dichloromethane. The majority of the compounds studied were of analytical grade (Aldrich or Ega, Merck, Fluka). All were analyzed by gas chromatography (GC). Those of purity greater than 99% were only dried by an appropriate reagent. The others were purified by distillation (on a spinning band column) or by preparative GC. Reagent grade sulfolane (Merck, 97%) was carefully crystallized three times and finally checked by GC on O.V. 17 (final purity > 99.9%). (9) (a) Preliminary communications: Elegant, L.; Fratini, G.; Gal, J. F.; Maria, P. C. Presented at the 'OnziEmes Journdes de Calorimttrie et d'Analyse Thermique", Barcelona, Spain, June 4-6, 1980 (Vol. 11, Abstract 3/21). (b) Maria, P. C.; Gal, J. F.; Elegant, L.; Azzaro, M. Presented at the 'Second Euchem Conference on Correlation Analysis in Organic Chemistry", Hull, England, July 18-23, 1982 (Abstracts p 39). (10) Spectroscopic grade dichloromethane is stabilized by 20 ppm of 2methyl-2-butene. Reagent grade CHIClz purified by treatment with concentrated H2SO4,I1distilled from P2O5l2and also stored in dark bottles over a 4 A molecular sieve gave identical calorimetric and spectroscopic results.
0 1985 American Chemical Society
A Lewis Basicity Scale for Nonprotogenic Solvents Calorimetry. The solution of the basic compound to be studied is prepared under an atmosphere of ultradry argon by weighing and volumetry, then dried on a 4-A molecular sieve for at least 4 h. The lower practical concentrations (for sparingly soluble compounds) are 0.2 to 0.3 mol.L-', taking into account the number of measurements possible on 3 cm3 of solution. Concentrations up to 1 mo1.L-' are used in some cases (solvents of low basicity) to prevent dissociation of the 1:l adduct. Within the abovementioned concentration range no significant differences in enthalpy changes were observed. The differential heat-flux calorimeter used in this study is of the Tian-Calvet type.I3 A borosilicate glass cell was dried at 393 K, flushed with ultradry argon, and inserted in the calorimeter. A capillary tube (2 mm i.d., 8 mm o.d.), ended by a sintered glass disk is plunged into 50 g of pure dry mercury14 which fills the bottom end of the cell. This tube is connected to a vacuum line used for BF3 storage and handling and evacuated to less than lo4 mbar. The solution (3 cm3) was poured into the cell (over the mercury). The cell, its content, and the "heat sink" of the calorimeter reach thermal equilibrium in less than 1 h (or about 20 time constants). When dichloromethane is used as a solvent, measurements are started only after 14 h (overnight) to allow CH2C12to reach the liquid-vapor equilibrium, thereby preventing evaporation and the resulting base line drift. A mercury manometer fitting in the cell allows the pressure on the solution to be checked. This experimental setup is a scaled down version of the Brown ca10rimeter.I~ The principle of our calorimetric procedure has been givenI6 and a more complete account of its current applications will be published elsewhere. We limit the description to the essential features as a means of evaluation of the accuracy of our method by contrast to some previous calorimetric works in this field. We operate by discontinuous titrations of the solute. Aliquots of about lo4 mol of gaseous boron trifluoride, from a calibrated volume, are added to the solution, the gas being permitted to pass under its own pressure through the sintered glass disk covered with mercury, acting as an effective one-way valve." Each addition of a quantity of BF3 ( n mol) generates a quantity of heat (Q in joules). n is the difference between the initial and the unreacted quantity of BF3 (P,V,T measurements). Q is obtained by integration (2105 M integrator for calorimetry and DTA, Delsi, Suresnes, France) of the amplified (A 31 amplifier, Setaram, Lyon, France) calorimeter signal, which is proportional to the heat flux dQ/dt. An absolute calibration is made by the Joule effect using a constant current power supply (E.J.P. 30, Setaram). The signal from the calorimeter is simultaneously recorded as a function of time (Servotrace, Sefram, Paris, France). The very fast complexation reaction appears as a sharp peak with a fast rise and a slower exponential ( t l 1 2 150 s) return to the base line. The maximum ordinate Y,,, of the recorded curve is in this case proportional to the instantaneous heat Qfas1'3a evolved. If any slow thermal phenomenon (Qslow) occurs after the fast reaction, Qfast Qslowis no longer proportional to Y,,,, so we can detect the presence of side reactions (polymerization, crystallization, etc.) much slower, in general, than the complexation. The corresponding measurements are rejected. The experimental enthalpy of complexation for each addition, AHo, is defined as the QfaSt/n
+
( 1 1 ) Riddick, J. A.; Bunger, W. B. "Techniques of Chemistry", Weissberger, A. Ed.;Wiley-Interscience: New York, 1970, Vol. 2, 3rd ed, pp
77C-77 1. (12) Gal, J. F.; Morris, D. G. J. Chem. Sac., Perkin Trans. 2 1978,
431-435. (13) (a) Calvet, E.; Prat, H. 'Microcalorim6trie, Applications PhysicoChimiques et Biologiques"; Masson: Paris, 1956. (b) Evans, W. J. 'Biochemical Microcalorimetry, The Conduction-Type Microcalorimeter", Brown, H. D., Ed.; Academic Prcss: New York, 1969;Chapter XIV. (c) McGlashan, M. L. "Chemical Thermodynamics";Academic Press: London,
1979. (14) Wilkinson, M. C. Chem. Reu. 1972,72, 575-625;see also Gal, J. F.; Azzaro, M. J. Chem. Educ. 1976,53, 118. (15) Brown, H. C.; Gintis, D. J . Am. Chem. Sac. 1956,78,5378-5383. (16)Gal, J. F.;Calleri, C.; Elegant, L.; Azzaro, M. Bull. SOC.Chim. Fr. I1979,311-319. (17)Brown, H.C.; Domash, L. J . Am. Chem. Sm. 1956,78,5384-5386.
The Journal of Physical Chemistry, Vol. 89, No. 7, 1985 1297 ratio. Our titration method allows us to obtain 6 to 8 AHo values using 3 cm3of solution. One or two titrations are usually sufficient to obtain good reproducibility on 8 to 12 AHo values. In fact, we observed in several cases that during a titration the first and the last -AH'O are too high and too low, respectively. Low values are attributed to a partial dissociation of the weak complexes.I6 Brown et al. have rejected the value corresponding to the first injection" as species more basic than the studied solute, though present in low, or very low, quantities, bring a contribution to the enthalpy change measured. The 1:l stoichiometry of the complex formed is controlled by accurate measurements of the BF3/donor ratio using two independent methods: one is the evaluation of the BF3 in excess by measuring its overpressure on the solution with the mercury manometer fitted to the calorimeter cell (free BF3 is scarcely soluble in CH2C12 a t 298 K). The other method consists of estimating the quantity qOt of complexed BF3 as QtOt/AHowhere Qtot is the total heat evolved during a titration. The structures of complexes were carefully checked by 'H N M R and IR spectroscopies.I2 Results. The enthalpy changes for reaction 1 for 75 basic, nonprotogenic solvents, as solutes in dichloromethane solution, are shown in Table I. The choice of this measuring medium will be mentioned again in the Discussion, but some arguments of an experimental nature are presented here: most boron trifluoride molecular complexes are only sparingly soluble in nondipolar solvents, so the choice was reduced to some aromatic or chlorinated solvents (CC14 excluded). Dichloromethane was chosen as the working medium owing to its low toxicity and its frequent use in spectroscopic studies of complexes. Moreover, it is weaker as a hydrogen bond donor than trichloromethane. The most important problem in the calorimetric work concerns its volatility but this was circumvented by working under dichloromethane equilibrium vapor pressure. Nitrobenzene also exhibits good solvent properties toward boron halides complexes.'* Its weak basicity and high dipolarity as compared to the weak acidity and low dipolarity of dichloromethane led us to run some measurements on representative bases in nitrobenzene (Table 11). A good linear relationship is observed between the enthalpy changes in dichloromethane and nitrobenzene (kJemo1-I): -Affo~~,(CH2c12 Soh) = 0.958[-hH0B~,(C6HSN02 soln)] - 0.31 (2) number of measurements: n = 12i correlation coefficient: r = 0.9970; standard deviation: s = 2.26 (established for compounds 5, 31, 36, 45, 48, 50, 55, 62, 68, 72, 101, and 102). With the intent to minimize the dissociation of weakly bound complexes, some measurements have been done on pure bases (bulk liquids). -AHoBF,(C6H5NO2bulk) (Table 11) allows us to calculate (eq 2) -AHoBF,(CH2C12soln) shown in Table I, as bulk nitrobenzene plays the role of solute and solvent. A series of experiments done on bulk nitromethane gave -AHoBF,(MeN02 bulk) = 39.60 kJ-mol-I; assuming that bulk nitromethane may be considered as nitromethane solubilized in nitrobenzene, similia similibus sol~untur,~ we calculate from eq 2 the corrected value reported in Table I. Di-tert-Butyl ketone was also investigated as a bulk liquid (-AhHoBF,(bulk) = 33.01 kJ-mol-') and corrected in the same way. For these three weak bases the corrections are very small. It is noteworthy that the regression line (eq 2) passes (within experimental error) through a point having as coordinates the calculated enthalpies of solution of BF3 in dichloromethane and nitrobenzene2' if these solvents act only by nonspecific effects. (18) Brown, H. C. J . Chem. Sac. 1956,1248-1268. (19)Woolhouse, R. A,; Eastham, A. M. J . Chem. Sac. B 1966,33-36. (20) (a) Derrieu, G.; Gal, J. F.; Elegant, L.; Azzaro, M. C. R . Acad. Sci. Paris, Ser. C 1974,705-707. (b) Azzaro, M.; Derrieu, G.; Elegant, L.; Gal, J. F. J . Org. Chem. 1975,40, 3155-3157. (21)Azzaro, M.; Gal, J. F.; Geribaldi, S.; Grec-Luciano, A,; Calleri, C. J . Chem. Res. ( S ) 1979,134-135. (22)Azzaro, M.; Gal, J. F.; Geribaldi, S. J . Chem. Sac., Perkin Trans. 2 1984,771-774.
1298 The Journal of Physical Chemistry, Vol. 89, No. 7, 1985
Maria and G a l
TABLE I: Enthalpies of Complex Formation with Boron Trifluoride in Dichloromethane for Nonprotogenic Solvents no. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38
solvent dichloromethane di-tert-butyl ketone nitrobenzene nitromethane dimethyl sulfite tetramethylene sulfone diethyl sulfite benzonitrile benzyl cyanide methyl benzoate isobutyronitrile acetonitrile tert-butyl cyanide propionitrile cyclohexyl cyanide n-butyronitrile ethyl benzoate propylene carbonate dimethyl carbonate diisopropyl ketone acetaldehyde methyl formate cyclobutanone di-n-butyl ketone diethyl carbonate ethyl formate diethyl ketone methyl acetate tert-butyl methyl ketone di-n-propyl ketone dioxane acetophenone isopropyl methyl ketone benzaldehyde ethyl acetate acetone ethyl methyl ketone propyl methyl ketone
- A H o ~ ~ , ( 2 9 815. K)/ kJmol-' 10.0 f 3.0b 31.32 f O.4lc 35.79 f 1.4OC 37.63 f 0.56c 51.27 f 0.46 51.32 f 0.29 55.13 f 0.77 55.44 f 0.28 56.61 f 0.23 59.4 f 1.ld 60.07 f 0.31 60.39 f 0.46 60.92 f 0.13 60.95 f 0.21 61.16 f 0.33 61.18 f 0.28 61.2 f O.Sd 64.19 f 0.39 67.63 f 0.38 68.07 f 0.74e 69.57 f 1.23/ 69.76 f 0.1 1 70.30 f 0.16 70.70 f 0.57' 71.03 f 0.35 71.17 f 0.29 72.28 f 0.23 72.79 f 0.33 72.83 f 0.37e 73.28 f 0.49' 74.09 f 0.27 74.52 f 0.158 74.84 f 0.25 74.88 f l.OOh 75.55 f 0.31 76.03 f 0.21 76.07 f 0.33 76.19 f 0.37
no. 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76
solvent cyclohexanone diisopropyl ether camphor cyclopentanone cycloheptanone di-n-butyl ether diethyl ether di-n-propyl ether trimethyl phosphate tetrahydropyran oxepane tetrahydrofuran isophorone 2,6-dimethylpyridine dimethylethyleneurea (DMEU) 2,4,6-trimethylpyridine dimethyl sulfoxide di-n-butyl sulfoxide tetramethylurea N,N-dimeth ylaniline N,N-dimethylformamide
N-isopropyl-2-pyrrolidone dimethylpropyleneurea (DMPU) N,N-dimethylacetamide N-meth ylpyrrolidone N,N-dieth ylformamide N,N-dieth ylacetamide 1-formylpiperidine N-meth ylpyridone hexameth ylphosphoramide tripiperidinophosphine oxide tripyrrolidinophosphine oxide 2-meth ylpyridine pyridine 3-methylpyridine 4-methylpyridine triethylamine N-meth y lpyrrolidine
-Aff0BF,(298.15 K ) / kJ.mo1-I 76.36 f 0.82 76.61 f 0.39 77.30 f 0.20 77.44 f 0.45 77.57 f 0.33 78.57 f 0.39 78.77 f 0.38 79.42 f 0.27 84.75 f 0.22 85.36 f 0.46 87.78 f 0.38 90.40 f 0.28 90.56 f 0.41' 97.73 f 0.58 98.93 f 0.38 101.03 f 0.29 105.34 f 0.36 107.60 f 0.46 108.62 f 0.22 109.16 f 0.76 110.49 f 0.18 112.08 f 0.99 112.13 f 0.29 112.14 f 0.41 112.56 f 0.36 113.20 f 0.35 113.61 f 0.25 114.16 f 0.57 116.94 f 0.33 117.53 f 0.45 118.14 f 0.92 122.52 f 0.14 123.44 f 0.47 128.08 f 0.50 130.93 f 0.39 134.17 f 0.59 135.87 f 1.67' 139.51 f 0.77
"Quoted errors are confidence limits at the 95% level; number of measurements ranges from 8 to 12. bEnthalpy change for solution of gaseous BF, in CH2C12,from the relation AHo,, = -AHo,, Affomix(see ref 16 and references therein). This result is in good agreement with the -10.24 kJ-mol-l value estimated from the solubilities oPBF, in CH2CI2at 193 and 183 K reported by Woolhouse and E a ~ t h a m . ' cCorrected ~ values from measurements on bulk liquids, see text. From a simultaneous determination of the equilibrium constant and the enthalpy change.I6 e Corrected values from ref 20, after improvement of the calibration technique. fThe dichloromethane solution of acetaldehyde (Fluka puriss from a nonpreviously opened bottle) not dried on a molecular sieve, to prevent catalytic oxidation, is investigated during the first few hours following its preparation. gFrom ref 21. *From ref 22. 'From ref 23. jSee text.
+
TABLE 11: Enthalpies of Complex Formation" with Boron Trifluoride in Nitrobenzene and Comparison with Literature Values no. 3 5 36 31 45 48 50 55 62 101 68 72 102
solvent nitrobenzene dimethyl sulfite acetone dioxane diethyl ether tetrahydropyran tetrahedrofuran dimethyl sulfoxide N,N-dimeth ylacetamide thiazole hexameth ylphosphoramide pyridine 4-(dimethy1amino)pyridine
-AHo~F,(C6H5N02Soh) 37.68 f 1.4OC 56.25 f 0.82 78.06 f 0.25 79.78 f 0.51 81.35 f 0.62 86.45 f 0.54 92.97 f 0.30 113.16 f 0.32 117.01 f 0.65 122.13 f 0.77 123.11 f 0.46 137.86 f 0.71 156.67 f 0.85
-hHo~~,(CH2C12
-AH0,~?(C6H5NO2 Soh) lit. values 38.66 f 0.50d
84.3' 86.82 f 0.71'
118.35 f 0.49g
126.65 f 0.57h 137.53 f 0.67f
151.55 f 0.76'
"All values in kJmol-l at 298.15 K. bSee table I unless otherwise noted. 'Experiment done with bulk nitrobenzene (considered as nitrobenzene solubilized in nitrobenzene). dFrom Henry's law constants: see ref 24. 'From ref 24 and 25. /From ref 15. ZThis work; not a usual solvent, but the existing literature value led us to investigate this compound. From ref 26, the author taking as standard an enthalpy change of -139.45 kJmol-' for pyridine in nitrobenzene. Correction using accepted value for 72 leads to a better agreement with our measurement. 'This work; solid base having the highest enthalpy change of complexation with BFj in CH2CI2and in C 6 H 5 N 0 2yet measured. T h e s e enthalpies a r e calculated from the regular solution theory.'6,28 Triethylamine reacts slowly with dichloromethane as
mentioned in t h e literature for several tertiary amine^.*^*^^ Few measurements a r e possible immediately after preparation of t h e
(23) Azzaro, M.; Gal, J. F.; Geribaldi, S . J. Org. Chem. 1982, 47, 4981-4984.
(24) Brown, H. C.; Holmes, R. R. J . Am. Chem. SOC.1956, 78, 2 173-21 76.
A Lewis Basicity Scale for Nonprotogenic Solvents solution. The reduced calorimetric stabilization time and the side reactions induce a decrease of accuracy and a slight overevaluation of its -APBF3(CH2C12 soh) (about 142 kJ-mol-'). Consequently, the value shown in Table I for 75 was calculated from -AHoBF,(C6H5N02s o h ) = 142.15 f 1.09 kJ.mol-' and eq 2.
Discussion Medium Effects. Drago and co-workers, in their search for a quantitative evaluation and prediction of donor-acceptor interactions based on the E and C model,5 were hampered by the lack of reliable data in the so-called "poorly solvating" media,3 i.e., carbon tetrachloride and hexane, chosen as the best substitutes for solvation free gas-phase medium. They proposed correction of the enthalpy changes measured in polar, weakly basic solvents, using a method termed ESP32(elimination of solvation procedure), which was applied to acid-base interactions involving various HBD and Lewis acids. ESP was later applied to the weakly acidic solvent CH2C12.30*33 In particular Drago et al. discouraged the use of this solvent for thermodynamic studies of the interaction between a reference Lewis acid such as iodine or BF3 and a series of basesem The ESP correction fails in this case because of a supposed hydrogen bonding between the solvent and the fluorine or iodine atoms in the adducts, the strength of this interaction being a function of the donor basicity. For BF3 adducts such hydrogen bonding, if present, should be very weak if one considers the mutual deactivating effect of the three f l ~ o r i n e s . In ~ ~this respect, the +BF3 moiety may be compared to a -CF3 (or -SF5) substituent, which was shown to be inert toward HBD solvents.34 Our aim is not to establish absolute (solvation free) bond strengths but rather a relative basicity scale. With this in mind, even if hydrogen bonding between dichloromethane and solutes and/or their BF3 adduct exists, it is sufficient that the energy involved be weak and roughly proportional to the solute's Lewis basicity. In our opinion, this is the case for the BF3/solute/CH2C1, system. This is supported by the linear relationship between - A P B F , measured in CH2Clzand in C6H5N02(eq 2). It should be noted that this relation covers a wide range of reactivity (about 100 kJ from dimethyl sulfite to 4-(dimethylamino)pyridine, see footnote i, Table 11). Literature values for the gas-phase complexation enthalpies, -AHoBF (gas), between boron trifluoride and diethyl ether,35ethyl acetate,js tetrahydropyran, and t e t r a h y d r o f ~ r a nshow ~ ~ a trend (25) Brown, H. C.; Horowitz, R. H. J . Am. Chem. SOC. 1955, 77, 1730-1733. (26) Azzaro, M. Bull. SOC.Chim. Fr. 1964, 2201-2202. (27) Ordinate: - A H o B F , = 10.0 f 3.0 kJ-mol-) for CH2CI2(footnote b, Table I); abscissa: -WBF, = 9.8 & 3.0 kJ-mol-' for C6HSNO2(solubility parameter 6 = 20.5 J0.5cm+ for C6H5NOzfrom ref 28). If this point is included in the correlation (the regression line is, of course, quasi-force-fitted through this calculated point distant from the experimental ones) we obtain - W B F , ( C H ~ C Soh) I~ = O . ~ ~ ~ [ - W B F ~ ( C Soh)] ~ H +S 0.07 N~~
n = 13, r = 0.9985, and s = 2.07 giving no significantly different values for the corrected -AHoBFshown for compounds 2, 3, 4 and 75 in Table I. (28) Barton, A. F. Chem. Rev. 1975, 75, 731-753. (29) Mills, J. E.; Maryanoff, C. A.; Cosgrove, R. M.; Scott, L.; Mc Comsey, D. F. Presented at the 186th National meeting of the American Chemical Society, Washington DC, Aug 28Sept 2, 1983, Division of Organic Chemistry, and personal Communication from Dr. Mills. (30) Drago, R. S.; Nusz, J. A,; Courtright, R. C. J. Am. Chem. Soc. 1974, 96, 2082-2086. (31) Drago, R. S. Struct. Bonding (Berlin) 1973, 15, 73-139. (32) (a) Nozari, M. S.; Drago, R. S. J. Am. Chem. SOC.1972, 94, 6877-6883. (b) Drago, R. S.; Nozari, M. S.; Vogel, G. C. J. Am. Chem. Soc. 1972, 94, 90-94. (33) Nozari, M. S.; Jensen, C. D.; Drago, R. S. J. Am. Chem. SOC.1973, 95, 3162-3165. (34) (a) Taft, R. W. personal communication. (b) Chawla, B.; Pollack, S. K.; Lebrilla, C. B.; Kamlet, M. J.; Taft, R W. J. Am. Chem. SOC.1981, 103, 6924-6930. (35) Mc Laughlin, D. E.; Tamres, M. J. Am. Chem. SOC.1960, 82, 5618-5621. (36) Sacks, L. J.; Drago, R. S.; Eyman, D. P. Inorg. Chem. 1968, 7 , 1484-1488. (37) Mc Laughlin, D. E.; Tamres, M.;Searles, Jr., S. J. Am. Chem. SOC. 1960,82, 5621-5625.
k.
The Journal of Physical Chemistry, Vol. 89, No. 7, 1985 1299 TABLE I11 Literature Complexation Enthalpies of Basic Solvents with SbCL in 1.2-Dichloroethane"
no.
2 3
-AH0sbc15(298.1SK)/kJ-mol-l 40.33 f 0.29b 34 f 3b
4 6 8 9 11 12 14 16 18 19 20 25 26 27 28 29 31 32 33
35 36 37 38 39 45 46 47 50 55 57 59 62 63 64 65 68 69 70 72 75
62.4 f 0.3b 63.47 f 0.13b 67.78 f 0.13b 66.86 f 0.21b 71.13 f 0.29b 69.16 f 0.13b 68.53 f 0.17b 70.92 f 0.13b
18.4c 1 1 .3c 619 49.8' 63.T 64.4c 59.OC 67.4' 693 63.2'
54.4 f 0.8' 61.1 f 0.4c
69.OC 619
33.99 71.42 f 0.13b 71.46 f 0.21b 71.25 f 0.17b 72.93 0.13b 73.22 f 0.08b 74.43 f 0.04b
69.0' 71.1c
69.5 f 1.3'
80.3' 74.64 f O.O@
88.0 f 0.2b 124.01 & 0.13b 116.32 f 0.13b
142.67 f 1.67' 132.6 f 2.5'
962 83.7c 124.7c 130d 113.3c 116.3c 1142 129.3' 134.7c 162.3c
138.5. 255d
85.8 f 2.1' 149XC 118.0h 21OSh 222h 197.5' 142.3 f 1.7' 202.5h
a Values considered reliable are italicized. References 6b and 40; reported errors are standard deviations. CGutmann's DN, ref 41; no reported experimental errors; some values of unknown experimental origin. d D N determined by an indirect method; ref 4b. 'Reference 6c. 'Reference 42. #Reference 43 (in benzene solution). Reference 44. 'Reference 45 (in dichloromethane solution). 'Reference 6a (in dichloromethane solution at 218 K).
very similar to those observed on - A H o B p 3 measured in CH2C12. -moBF3(gas) are available for other bases and preliminary results of a comparative study have shown us that, on a greater range of reactivity and for compounds widely different in their basic site, this trend is confirmed. These observations on medium effects lead us to think that we have reached one of the main goals of this work, namely a solvent Lewis basicity scale free from large and unpredictable solvation effects. Critical Examination of the DN Scale. Gutmann presented the D N scale as a general basicity scale and many workers have followed him using this parameter in that sense.38 He attempted to prove that basicity order remains constant even when using reference acids other than SbClSs4This point was disputed by (38) Among the numerous papers using DN parameters at face value and without any reservation about their applicability to solvent effects, see for example, the following: General treatment: (a) Krygowski, T. M.; Fawcett; W. R. J. Am. Chem. SOC.1915, 97, 2143-2148. Principal component analysis: (b) Fawcett, W. R.; Krygowski, T. M. Can. J. Chem. 1976, 54, 3283-3292. (c) Chastrette, M. Tetrahedron 1979, 35, 1441-1448. Phase transfer catalysis: Antoine, J. P.; de Aguirre, I.; Janssens, F.; Thyrion, F. Bull. SOC.Chim. Fr. 2 1980, 207-232. Dissociation of homoconjugated aliphatic carboxylate anions: Pawelka, Z.; Haulait-Pirson, M. C. J . Phys. Chem. 1981, 85, 1052-1057.
1300 The Journal of Physical Chemistry, Vol. 89, No. 7, 1985 6.
. ’ 72
65. 6.
55
nitrobenzene6b (Table 111) is close to our -AHoBF3. Notice also that Gutmann gives very different D N values for nitrobenzene and nitromethane although close basicity are expected on structural grounds, and this is what we observe (Table I). At the other end of Gutmann’s scale the incredibly overevaluated value for triethylamine of unknown experimental source was already criticized by Taft et a1.6a Another very important solvent also presents an overevaluated DN: HMPA (68)departs significantly from the general trend observed in Figure 1. Impressive is the near 50 kJ difference between the values reported by Gutmann and by Bollinger et al., showing that serious problems arise in measuring the basicity of this compound toward SbCIS. To try to understand the possible reasons for this anomaly we return to the definition of D N and to its measurements. DN has been defined as the negative AH values for 1:1 adduct formation between antimony pentachloride and electron pair donor solvents D, both in dilute 1,tdichloroethane solution according to
.
75
0
*
.
61
63
-59
e7
.
Maria and Gal
10
‘5.
D(so1n)
+ SbCl5(soln) s D:SbC15(soln)
(3)
AHosw15 for this reaction is the difference between the two terms AH1 and AH2, AH1 corresponding to 15
50
75
100
115
110
-Abli~~/kJm . ol-’
Figure 1. Plot of -AHosbc15 vs. -AHoBF3; numbering as in Table I; data points correspond to the italicized values in Table I11 (somepoints and numbers are omitted for clarity); data point 68 refers to Gutmann’s DN for HMPA.
Drag0~~9~’ who stated in particular that linear relationships between different basicity scales can be obtained if the EB/CB ratio is constant. For example, a detailed study of carbonyl compounds basicity toward iodine shows such relationships because bases have an almost constant EB/CB ratio.39 Another condition for obtaining proportional basicity scales is that the reference acids have the same EJCA ratio (about 1.5 in the case of SbC15). In addition to electronic effects, steric strain18 must be taken into account in the studies of Lewis acid-base interactions. Now the user, fully aware of the application range of DN, meets with an unexpected obstacle: some odd DN values, probably due, in part to the calorimetric method used, and also in part to the nature of SbCl,. Enthalpy changes for the SbCls adduct formation in 1,2-dichloroethane have been reported by several authors and literature data on compounds studied here are collected in Table 111. Comparison with BF3 complexation (Table I) shows that thermic effects are roughly equal for the same base. A plot of -Wsbcl, vs. -WBF, for the italicized values in Table I11 (Figure 1) shows evidence of a linear trend with a slope of about 1. A better fit should be expected owing to the similar strengths and steric requirements of BF3 and SEI5. (Compare, for example, -WBF, and - A H O S K I ~ (italicized) for compounds 2, 12, 20, 45, 50, 75). However, several of Gutmann’s D N values are not in agreement with this statement. The DN for weak bases reported by this author are underevaluated, as noted by Kolling,46because no correction was made to take into account the dissociation of the By contrast, the value reported by Olofsson for (39) Laurence, C.; Guiheneuf, G.; Wojtkowiak, B. J . Am. Chem. SOC. 1979. 101. 4793-4801. (40) (a) Olofsson, G.; Lindquist, I.; Sunner, S. Acra Chem. Scand. 1963, 17,259-265. (b) Olofsson,G. Acta Chem. Scand. 1964,18, 11-17. (c) Ibid. 1964, 18, 1022-1023. (d) Zbid. 1965, 19, 2155-2159. (e) Ibid. 1967, 21, 2143-2150. (4Ibid. 1967,21,2415-2422. (g) Ibid. 1968,22, 1352-1353. (41) (a) Gutmann, V.; Steininger,A,; Wychera, E. Monatsh. Chem. 1966, 97, 460-467. (b) Gutmann, V.; Wychera, E.; Mairinger, F. Ibid. 1966, 97, 1265-1275. (c) Gutmann, V.; Mayer, U. Ibid. 1967, 98, 294-297. (d) Gutmann, V.; Scherhaufer, A. Ibid. 1968, 99, 335-339. (42) Arnett, E. M.; Petro, C. J. Am. Chem. SOC.1978,100, 5402-5407. (43) Paul, R. C.; Ahluwalia, S . C.; Parkash, R. Indian J . Chem. 1968,6, 464-465. (44) Bollinger, J. C.; Yvernault, G.; Yvernault, T. Thermochim. Acta 1983, 60, 137-147. (45) Ozari, Y.; Jagur-Grodzinski,J. J. Chem. Soc., Chem. Commun. 1974, 295-296. (46) Kolling, 0. W. Anal. Chem. 1982, 54, 260-264.
D( 1)
+ SbC15(soln) G D:SbCIS(soln)
(4)
and AH2 corresponding to D(l)
F!
D(so1n)
(5)
The former AH1 is measured in a suitable ~ a l o r i m e t e rby ~~ breaking a sealed glass ampule, containing the donor, in a solution of an excess of the Lewis acid. Impurities accompanying the donor are also complexed increasing or lessening AHl accordingly as they are more or less basic than D; in particular, the role played by water was well studied by O l o f ~ s o n . ~This ~ experimental procedure does not allow control of the stoichiometry of the complex formed in a medium containing an excess of acid. We have observed that strong donors can accept up to 1.1 mol of BF3 per mole of base in the presence of an excess (overpressure) of acid. It should be noted that our W Bcorrespond F3 to a titration region where there is an excess of the base. Gutmann et al. have operated by thermometry in an “adiabatic” calorimeter which acts as an integrator of the various heat sources and thus cannot reveal side reactions. In comparison, these drawbacks are avoided by using our isothermal heat-flux method which allows the monitoring of the kinetics of the reactions as emphasized in the Experimental Section. Moreover, with our vacuum line apparatus, the moisture-sensitive Lewis acid is never in contact with the atmosphere. It should be noted also that using pure bases (reactions 4 and 5) led to measurements of solution enthalpies AH2 generally small, exo- or endothermic (with the associated calibration problems). Finally, SbCIS is known as an oxidizing49and chlorinatingSo reagent. Conductivity measurements show ions formation in the presence of acet~nitrile,~~ tertiary phosphine^,^^ acetic anhydride,s3 and other donors.& It is of particular interest to discuss the cases of triethylamine and HMPA of dubious DN values. Bollinger et al.44 found that triethylamine is rapidly quaternized by 1,2dichloroethane under the catalytic action of SbCISas suggested earlier by Taft et a1.;6a in similar concentrations (2.2 X mo1.L-I) 1,2-dichloroethane solutions of SbC15 + Et3N and of (47) Gutmann, V.; Mairinger, F.; Winkler, H. Monatsh. Chem. 1965, 96, 574-580. (48) Olofsson, G. Acta Chem. Scand. 1967, 21, 1887-1891. (49) Smith, J. D. In “Comprehensive Inorganic Chemistry”, TrotmanDickenson, A. F., Ed.;Pergamon Press: Oxford, 1973; Chapter 21, p 547. (50) Cotton, F. A.; Wilkinson, G. ‘Advanced Inorganic Chemistry”, Interscience-Wiley: New York, 1962. (51) Kolditz, L. In “Advances in Inorganic Chemistry and
Radiochemistry”,Emeleus, H. J., Sharpe, A. G., Ed., Academic Press: New York, 1965; Vol. 7. (52) Ruff, J. K. Inorg. Chem. 1963, 2, 813-817. (53) Viard, B.; Poulain, M.; Grandjean, D.; Amaudrut, J. J . Chem. Res. (S) 1983, 84-85. J . Chem. Res. (M) 1983, 853-867.
The Journal of Physical Chemistry, Vol. 89, No. 7, 1985 1301
A Lewis Basicity Scale for Nonprotogenic Solvents
TABLE I V Correlations between Various Basicity Dependent Properties (BDP) and - A H o B F , by Use of the Model BDP = a ( - A H o B F 3 ) BDP
BDPn
a
-AHosbc15/kJ.mol-’ - ~ o , , / k l . m o l - l8 B(AvoH)/cm-l’ B( A v ~ o/Cm-’ )
-4.83 -14.32 -106.33 -59.44 0.061 4.99
1.09 0.37 4.42 2.38 0.0057 0.22
tyoe of oroperty Lewis basicity hydrogen bonding
8“
’
-AHopFp/kJ.mol-’
a,b
36 31h 39, 33’ 43“ 1 9p
re
Sd
0.9628 0.9021 0.9457 0.9525 0.8290 0.8218
7.42 4.55 37.79 19.8 1 0.095 3.80
+ BDPo
“Number of solvents included in the correlation. bSolvents supposed to present steric effects toward BF3 are excluded from data set others than Lewis basicity scales. Correlation coefficient. dStandard deviation. From Olofsson, Gutmann, and Drago, see Table 111 and related footnotes. /Only italicized values (Table 111) are considered in the correlation. gComplexation with iodine from ref 39, 57, and 58; measurements in heptane and tetrachloromethane were used. hSolvents 5, 6, 12, 17, 19, 25, 26, 28, 29, 31, 32, 33, 34, 35,36, 37, 45, 47, 48, 50, 55, 57, 58, 59, 62, 65, 71, 72, 73, 74, 75. ‘ B = Av$& = - ~$fAk...~ from ref 58 and 59. /Solvents 3, 4, 6, 8, 9, 10, 11, 12, 13, 14, 16, 17, 18, 25, 28, 31, 32, 34, 35, 36, 37, 39, 40, 44, 45, 46, 47, 48, 50, 55, 57, 58, 59, 62, 65, 68, 71, 72, 74. ’ E = A v M ~ D= v M a D - YM~D-B from ref 60 and 61. ‘Solvents 3, 4, 8, 12, 14, 20, 25, 27, 28, 29, 31, 32, 33, 35, 36, 37, 38, 39, 40, 42, 44, 45, 46, 50, 55, 58, 59, 62, 68, 72, 73, 74, 75. “Solvatochromic parameter of HBA basicities from ref 62. “Solvents 3, 7, 8, 10, 12, 14, 17, 18, 19, 22, 25, 26, 27, 28, 29, 33, 34, 35, 36, 37, 38, 39, 40, 42, 43, 45, 46, 47, 48, 50, 55, 56, 57, 59, 62, 63, 65, 67, 68, 72, 73, 74, 75. OEnthalpies of hydrogen bonding with p-fluorophenol (PFP) in CC14 from ref 56. PSolvents 5, 12, 18, 25, 31, 35, 36, 37, 39, 42, 45, 47, 50, 55, 58, 62, 68, 72, 75. TABLE V Dual Parameter Correlations between Bonding: -AHoaF. = a x + b Y + c X Y
and Indicators of Electrostatic and Covalent Characters of the Donor-Acceptor
A H O B ~ ~
Gd -AHopm/kJ.mol-’
‘
t g
-API,/kJ.mol-’
‘
a
b
C
72.32 108.84 1.97
8.19 59.29 1.69
-12.50 22.58 12.12
n“ 1 4c 35h
19
Rb
sc
0.9561 0.9603 0.9727
7.40 6.41 5.65
“Number of data points; bases for which steric effect is expected: 40, 44, 57, 75 (see Drago, ref 31) and ortho-substituted pyridines 52, 54 and 7163.71 are excluded. bMultiple correlation coefficient. CStandard deviation. CSolvents 12, 28, 31, 35, 36, 45, 48, 50, 55, 59, 62, 68, 72, 74. /Reference 62. gA coordinate covalency parameter from ref 62 and 71. hSolvents 7, 10, 17, 18, 19, 22, 25, 26, 27, 28, 29, 33, 34, 35, 36, 37, 38, 39, 42, 45, 46, 47, 48, 50, 55, 56, 59, 62, 63, 65, 67, 68, 72, 73, 74. ’See footnotes g and 0, Table IV. /Solvents 12, 25, 31, 35, 36, 37, 45, 47, 50, 55, 58, 59, 62, 72, 74.
Et,N+Cl- give about the same conductivity: 1.9 X lo4 and 1.6 X lo4 52-l cm-l, respectively. A conductivity of 0.26 X IO4 i2-l an-’exhibited by a 0.86 X IF2mo1-L-l 1,2-dichloroethanesolution of SbC15 HMPT,44together with the preceding and other literature v a l u e ~ , 6 C *suggests ~ ~ - ~ ~ that this solution, contrary to the authors’44assumptions, contains a nonnegligible amount of charged species. We think that ions formation brings a contribution to the total heat quantity measured, all the more in the presence of an excess of SbC15. Relationships between -AHoBF,and the Most Used Basicity Scales Relevant to Solvent Effects. In a review on the correlations among various solvent parameters Griffiths and PughS5pointed out that the most used basicity scales are fairly well related and used the D N vs. 0-D stretching frequency of CH30D to expand the D N scale. Reasonably good correlations between various acid-base solvation parameters were also observed by Arnett et but they put forth, in the conclusion of their paper, that the energies for acid-base interaction, from a general viewpoint, do not correlate very well with each other, even when the processes are quite similar. We have reexamined such correlations, in the context of solvent effects, taking advantage of our accurate, wide range, Lewis basicity scale. Results of correlation analysis are given in Table IV. A first inspection shows that there is no very good or even fair relationship between any of the currently used basicity scales and -AW’B,. The possible causes of the poorer-than-expected fit with the D N scale has been discussed (vide supra). As anticipated the ’soft” acid I2 behaves somewhat differently from the “hard” acid BF3, even toward our series of “hard” bases (within this data set small differential steric effects may account in part for the poor fit). We shall return later to this “soft” basicity scale. Among hydrogen bonding scales we observe that the two infrared B scales are roughly proportional to -AW‘BF3. Large deviations from the correlations reported in Table IV are observed for sterically hindered bases (see footnote b, table IV). The larger electronic demand exerted by BF3 than by hydroxylic HBD also induces
+
(54) Gutmann, V.; Imhof, J. Monatsh. Chem. 1970, 101, 1-10, (55) Griffiths, T. R.; Pugh, D. C. Coord. Chem. Rev. 1979.29, 129-211. (56) Amett, E. M.; Mitchell, E. J.; Murty, T. S. S. R. J . Am. Chem. SOC. 1974, 96, 3875-3891.
deviations, as we showed recently within a series of substituted ~yridines.6~Other problems may also arise from the possibility of multiple basic sites or from purely spectroscopic problems such as band overlaps and couplings.64 When limited to solvent effects the relative similarity of “hard” Lewis basicity (DN or -WBF3) and the spectroscopic infrared scales ( 8 s ) may explain why such parameters were used indifferently as secondary parameters in addition to acidity, dipolarity, polarizability, and cavity parameters in the poineering works on multiparameter correlation analysis of solvent effects. In addition to other parameters, Krygowski and F a w ~ e t t , ~ ~ Glikberg and Parker et a1.,66and M a ~ e used r ~ ~D N while Koppe161and Shorter@‘ preferred their own hydrogen bonding basicity scales (see Table IV). If the solvent basicity is only a secondary effect the success of these multiparametric equations may be explained by the rough equivalence of the solvent basicity scale used. The result obtained with the @ scale deserve special comment: the @ scale has the largest data set in common with our scale and @ is the parameter having the weakest tie with our Lewis basicity scale, even if it was reported that @ correlates with DN.& In fact, the problem that @ cannot correlate all basicity-dependent (57) Gur’yanova,E. N.; Gol’dshtein, I. P.; Romm, I. P. “Donor-Acceptor Bond”; Wiley: New York, 1975. (58) Laurence, C.; Queignec-Cabanetos,M.; Dziembowska, T.; Queignec, R.; Wojtkowiak, B. J . Am. Chem. SOC.1981,103, 2567-2573. (59) Koppel, I. A.; Paju, A. I. Org. Reacr. (Tarru) 1974, IZ, 121-136. (60) Burden, A. G.; Collier, G.; Shorter, J. J . Chem. SOC.,Perkin Trans. 2 1976, 1627-1632. (61) Koppel, I. A.; Palm, V. A. in ‘Advances in Linear Free Energy Relationships”, Chapman, N. B., Shorter, J. Ed.; Plenum Press: London, 1972: Chaoter 5. (62) dmlet, M. J.; Abboud, J. L.; Abraham, M. H.; Taft, R. W. J . Org. Chem. 1983, 48, 2877-2887. (63) Berthelot, M.; Gal, J. F.; Laurence, C.; Maria, P. C. J. Chim. Phys. Phys.-Chim. Biol. 1984, 81, 327-331. (64) Laurence, C., personal communication. (65) Glikberg, S.; Marcus, Y. J. Solution Chem. 1983, 12, 255-270. (66) Parker, A. J.; Mayer, U.;Schmid, R.; Gutmann, V. J. Org. Chem. 1978, 43, 1843-1854. (67) (a) Mayer, U.Monatsh. Chem. 1978,109,421-433. (b) Ibid. 1978, 109; 7ii-790. (68) (a) Burden, A. L.; Chapman, N . B.; Duggua, H. F.; Shorter, J. J . Chem. Soc., Perkin Trans. 2 1978, 296-303. (b) Aslam, M. H.; Collier, G.; Shorter, J. J . Chem. SOC.,Perkin Trans. 2 1981, 1572-1576.
Maria and Gal
1302 The Journal of Physical Chemistry, Vol. 89, No. 7, 1985
1.1
i
J
I .D
74
7 1.
im
.
.
I1
0.3
I bo
80
100
-AH&
110
140
1bO
80
b0
a(-AH&
/kJ.mol-’
/hJ
.m 0 l - l
100
) + b(-AH
1, / k J . m o l - ’
110
1.0
J
)tC
Figure 2. 6 plotted against -AHoBF3; values for bases others than common solvents will be reported e l ~ e w h e r e . ~ ~ * ’ ~
Figure 3. -ANoBp3 as a function of -AHopm and -Mol,; numbers for data points correspond to numbers listed in Table I.
properties has not escaped the attention of Taft and Kamlet and they have very recently shown that a two-parameter equation is necessary.62 Figure 2 shows a plot of (3 vs. -AHoBF3 9b where a family-dependent6, behavior is apparent. If we exclude bases having large steric hindrance to the approach of BF3, crudely parallel lines can be drawn when families of bases having similar types of basic sites are considered separately. This separation seems to correspond to a greater covalent vs. electrostatic bonding character in the interaction with BF, than in that with a hydrogen bond donor molecule. In principle, the electrostatic and covalent components in a property related to the basicity may be separated by means of a dual parameter equation,69which is a generalized form of the Drago e q ~ a t i o n . ~Applied ,~’ to -AHo,, this dissection gives a significant fit, though of relatively low precision (Table V). To explain the residual deviations some reasons may be invoked including the intervention of a steric effect with BF370aeven if the more hindered bases have been removed from the data set,70b or the biased set containing only solvent type bases having a nitrogen or oxygen atom as a coordinating site. Another approach to the covalent character of the coordinative bond was recently proposed by Kamlet et al.62371by introducing a coordinate covalency parameter [to be used in combination with 8. The [ parameter is taken to be a property of a whole family of bases, as suggested by the family-dependent behavior and the parallelism of the lines both shown in Figure 2. The 8, [ correlation (Table V) works similarly to Eg and CBbut on a larger data set. Both E g , Cg and 8, [ parameters are the results of experimental data treated by models which can smooth off small deviations arising in part from experimental errors but also from physical effects ignored by the models. To avoid this problem one can directly use empirical parameters such as enthalpies of complexation, an obvious choice for our purpose. Following Drago7&and Panchenko et phenol (or p-fluorophenol, PFP) and iodine may be taken as reference acids73 owing to the re-
TABLE VI: Relationships between Proton and Alkali Metal Cations Gibbs Energy of Transfer from Water to Nonprotogenic Solvents and - ~ O B F ;
-AGo,(Mt,H20-S) Mt H+e Lit Nat
Kt Rb+
CS’
a 0.99 0.88 0.49 0.43 0.36 0.37
b -91.32 -77.01 -42.15 -35.08 -29.94 -27.94
+
= a(-MoBF3) b nb 6 7 8 11 9 8
re 0.8908 0.9680 0.9642 0.9926 0.9683 0.9374
Sd
18.15 8.46 5.78 2.07 3.71 5.46
’Units: kJ.mol-I; AGO, on mol-L-I scale at 298.15 K from Marcus.’* Transfers to 1,l-dichloroethane and to dichloromethane are assumed to give the same AGO,. ”See corresponding footnote to Table IV. eLack of important solvents to obtain a significant relationship.
spective electrostatic and covalent characters involved in their bonding with bases. The results of the two-parameters correlation between - A Z - Z o g ~ , vs. -AHoPFP and -AHo12are shown in Table V and in Figure 3 and are, at least, as good as those obtained from Eg, CBor 8, [ treatment^.^^ It should be noted that -AHopFpand -AHo12are quasi-orthogonal (in the statistical sense) giving more weight to this correlation. Physical significance may thus be attributed with reasonable qnfidence to the regression coefficients. The intercept (12.12 kl.mol-’) agrees well with the -10.0 kJ-mol-’ BF, enthalpy of solution in dichloromethane corresponding to the absence of any basic compound in the respective measurements media. The slopes are dose, showing that BF3 exhibit a behavior half-way between PFP and I,, and their values indicate about twice a sensitivity to structural effects in the bases for BF3 compared to PFP or I,. So we think we have gone further in the description of the BF, “hard-soft” character as well as its relative “ ~ t r e n g t h ” . ~ ~ -MaBF, as a Measure of the Solvation of Alkali Metal Cations. Transport of alkali metal cations through membranes plays a central role in biological systems. In this regard unraveling the effects implicated in the complexation of these “guest” ions by “host” carriers (solvents, macrocycles, etc.) may be aided by studies of their interaction with various organic ligands. For this purpose thermodynamic functions of transfer of these monovalent
(69) (a) Doan, P. E.; Drago, R. S. J . Am. Chem. SOC.1982, 104, 45244529. (b) Ibid. 1984, 106, 2772-2774. (70) (a) Already noted by Drago et al.: Drago, R. S.; Vogel, G. C.; Needham, T. E. J . Am. Chem. SOC.1971, 93, 6014-6026. (71) Kamlet, M. J.; Gal, J. F.; Maria, P. C.; Taft, R. W. J. Chem. Soc., Perkin Trans. 2, in press. (72) Panchenko, B.; Oleinik, N.; Sadovsky, Yu.;Dadali, V.; Litvinenko, L. Org. React. (Tarfu) 1980, 17, 65-87. (74) Small residual steric effects may be included in Mol: Rafik, C.; (73) We have selected p-fluorophcnol (PFF’) enthalpies of complexation Abboud, J. L. M.;Guiheneuf, G. J. Org. Chem. 1983, 48, 4761-4763. covering a large range of basicity and all measured in the same l a b o r a t ~ r y . ~ ~ (75) Pearson, R. G. ‘Hard and Soft Acids and Bases”; Dowden, Hutchinson and Ross: Stroudsburg, PA., 1973. For iodine complexation data sources see footnote g, Table IV.
The Journal of Physical Chemistry, Vol. 89, No. 7, 1985 1303
A Lewis Basicity Scale for Nonprotogenic Solvents
TABLE VII: Slopes of the Equations - A G o , ( M + , H 2 h S ) = 8 (-AHogp3) + b Compared to the Cas-PhaseAttachment Enthalpies of Four Water Molecules to Proton and Alkali Metal Cations
M+ H+ Li+ Na+ K+ Rb'
cs+
slope4 0.99 (1.4S)
0.88 0.49 0.43 0.36 0.37
-AHo /kJ.mol-' 1000.0 405.8 309.6 242.3 221.8 199.2
From Table VI. bAttachment enthalpies of four water molecules. Sum of the enthalpy changes for the reactions Mt(H20)rl + H 2 0* M+(H20), with n = 1 to 4.85 AHo for H t is taken as the sum of the attachment enthalpies of three water molecules to H 3 0 t and the proton affinity (PA) of H20.85 CWith the strongly deviating solvent 3 excluded.
IO
ao
,
50
70
-"HiF,/
(0
110
130
kJ.mol-'
Figure 4. -AGo,(K+,H2ChS) plotted against -iwOBF,;number of data points is as in Table I.
ions from water to nonaqueous solvents are of great interest. Recent attempts to rationalize ions Gibbs energy of transfer from water to organic solvents, -AGo,(MrrC,H20-S) have been made and empirical equations, including D N as a basicity parameter, We have reconsidered these correlations, were in light of our Lewis basicity scale, using a consistent set of selected data proposed by Marcus?8a Significant correlations are observed for -AGo,(M+,H20-S) vs. - A H o B F as shown in Table VI for alkali metal cations7Eb(H+ is included for comparison purposes). The standard deviations are to be compared to the 5 kJ-mol-' variations between AGO, from various sources, considered by Marcus as the maximum admissible for proposing a reliable value. Several comments are in order and are best exemplified by looking a t Figure 4 which shows the behavior of the most thoroughly studied ion, namely, K+. HMPA (68)and nitro compounds (3, 4) of widely different basicities fit the linear relationship well. We observe that the regression line passes through the points for dichloromethane (see footnote a, Table VI) and acetone of low (e = 8.93) and medium (e = 20.70) dielectric constants, respectively, when all other solvents are of > 30. This signifies that the solvent reaction field (ion-dipole interactions) plays a secondary role in alkali metal cations solvation. We have found that some improvements may be obtained by use of a multiparameter treatment including ion-dipole interaction^^^ and cavityE0 parameters in addition to -AHoBF,.Nevertheless the major part of the variances of the dependent variables are explained by the solvent Lewis basicity and the residual variances (see the standard deviation in Table VI) are close to the expected uncertainties in -AGO,; so we do not want to conclude positively on the role of these secondary effects, with the data currently available. A clear trend appears in the order of sensitivities of the different cations to the Lewis basicity, Le., the slopess1 of the correlation (76) Gritzner, G. Inorg. Chim. Acta 1977, 24, 5-12. (77) Abraham, M. H.; Liszi, J.; Kristbf, E. Aus?. J . Chem. 1982, 35, 1273-1 279. (78) (a) Marcus, Y . Pure Appl. Chem. 1983, 55, 977-1021. (b) Duly
warned by the author we have rejected in our treatment values given as 'uncertain". Sulfolane was found to deviate in all the correlations and was also excluded. Nevertheless, some "uncertain" values fit our correlations well. Selected AGO, for Ag+ and TI+do not fit the relationship in Table VI owing to their 'soft" ~haracter.'~ (79) (6 - 1)/2c is the relevant function; see ref 3a. (80) 6* (squared Hildebrand solubility parameter),8' approximately equal to the cohesive energy density, is homogeneous with the other thermodynamic quantities. See ref 28 and for a more recent review: Barton, A. F. M. "Handbook of Solubility Parameters and Other Cohesion Parameters"; CRC Press: Boca Raton, FL, 1983.
(Table VI): Cs+ N Rb+ < K+ < Na+ < Li+ < H+. We assign this tendency to the increasing overlap of the first solvation shell ligand's electron pairs orbitals and the cation's empty orbitals. This overlap should be reflected in the affinities of these cations toward a reference base such as the water molecule. To represent the first solvation shell, the cation are s u p p e d to be coordinated to four ligands.s2 In Table VI1 we compare the gas-phase attachment enthalpies of four water molecules clustering the cations5 to the corresponding sensitivity toward basicity. Parallel variations are observed,&confirming that Gibbs energies of transfer of alkali metal cations are mainly governed by some sort of "intrinsic" acidc-basic interaction, at least in the case of nonprotogenic solvents. This coherent description of cation solvation leads us to propose our - A H o B F 3 scale as an evaluation tool to test the reliability of not-yet-well-established N M R of alkali metal ions has also been used as probe of the immediate environment of the corresponding monovalent cations. Popov et al. have reported N M R data for 'Li, 23Na,8739K,8Eand W s E 9salts in different solvents. Popov observed a good relationship between infinite dilution 23Na ion chemical shifts and D N in ten solvents; an extrapolation of the D N scale to very basic solvents (ammonia and primary amines; triethylamine is not mentioned) was proposed. For Li+, K+,and Cs', such nice correlations are not observed; therefore it is clear that effects other than solvent basicity are operating on these properties and therefore solvent-induced chemical shifts of these ions cannot be used as a general measure of solvent "donicity". Even if the sodium nucleus behaves surprisingly well, owing to some supposed peculiar magnetic proper tie^:^ extrapolated D N for ammonia and primary aminesg0are clearly overevaluated. This is due to the biased slope for the reported c o r r e l a t i ~ nas~a~consequence ~~ of odd DN values, principally for the extreme points MeNOz and HMPA of underevaluated and overevaluated Lewis basicity, respectively. When the 6(23Na) are plotted against -AFsbc15 values calculated from the equation -AHosbc15 = BDPo + a(-AHoBF3) (see Table IV) rather than against Gutmann's DN, extrapolated DN for ammonia and primary amines are found about 50 kJ lower. The linear relationships 6(23Na) vs. calculated -AHososbc15 or vs. - A F B F , are (81) Sensitivities to - M 0 e ~ ,are barely affected by multiparameter treatments. (82) This number is a fairly good a proximation of the mean solvation number in waters3 and other (83) Gordon, J. E. "The Organic Chemistry of Electrolyte Solutions"; Wiley: New York, 1975. (84) (a) Davidson, W. R.; Kebarle, P. J. Am. Chem. SOC.1976, 98, 6125-6133. (b) Covington, A. K.; Newman, K. E. Pure Appl. Chem. 1979, 51,2041-2058. (85) Kebarle, P. Annu. Rev. Phys. Chem. 1977,28, 445-476. (86) Changing the nature to the reference base ( H 2 0 replaced by NH,)
or the number of ligands does not affect the affinity order and the qualitative aspect of this observation. (87) (a) Erlich, R. H.; Roach, E.; Popov, A. I. J . Am. Chem. Soc. 1970, 92,4989-4990. (b) Erlich, R. H.; Popov, A. I. J . Am. Chem. Soc. 1971,93, 5620-5623. (c) Popov, A. I. Pure Appl. Chem. 1975, 41, 275-289, and references therein. (88) Shih, J. S.; P o w , A. I. Inorg. Nucl. Chem. Le??.1977, 23, 105-1 10. (89) De Witte, W. J.; Liu, L.; Mei, E.; Dye, J.; Popov, A. I. J. Solution Chem. 1977, 6, 337-348. (90) Herlem, M.; Popov, A. I. J . Am. Chem. SOC.1972,94, 1431-1434.
1304 The Journal of Physical Chemistry, Vol. 89, No. 7, 1985
not very satisfa~tory.~'A multiparameter treatment similar to those used for transfers of alkali metal cations does not improve the fit, showing that the unexplained variance arises from experimental errors and perhaps from effects not directly related to classical solvation models.
Conclusion The aim of this work was to demonstrate the specificity and the usefulness of a Lewis basicity scale with particular emphasis on solvent effects. We have shown that the widely used Gutmann D N scale suffers the lack of reliability, principally from experimental problems (reference solvent and Lewis acid reactivities, together with the calorimetric method). We have developed an original experimental method leading to accurate measurements and allowing the detection of possible side reactions. The stoichiometry of the complex is controlled and the adduct's molecular structure is checked by spectroscopy. Measurements taken comparatively in PhN02 and in CHzC12do not show large differential solvation effects, so we assume the entire set of -AHoBF, is free from nonregular effects. We think that our measurements constitute a comprehensive and reliable Lewis basicity scale for nonprotogenic solvents. In the field of solvent effects, the calorimetric scales (-WBF, and -AHosbcc)are related in some way to the spectroscopic scale (AvOH and AvoD). This accounts for the indifferent use previously made of B or D N in literature as complementary scales together with other parameters (acidity, dipolarity, polarizability, and cavity). The hydrogen bond basicity parameter 0 is not related to -eBF, by a single line but rather by a family-dependent correlation. Lewis basicity is more sensitive to steric effect than hydrogen bond basicity and such an effect cannot be taken into account owing to the lack of parameters measuring the steric hindrance to solvation. We have shown the existence of a relation between -AHoBF3 and -AHOpFp together with -AHo, leading to the conclusion that BF3 has an acceptor behavior half-way between PFP and I2 and an acidic strength twice as large as those of PFP and 12. The - e B F , scale presented here appears as a usefull tool for the rationalization of the Gibbs energies of transfer of the alkali Summary and
(91) 6(23Na) vs. -MHOBF,: r = 0.8935 for the ten solvents used by Popv et
Maria and Gal metal cations which depend mainly on the solvent Lewis basicity. The decrease in ion solvation accompanying the increase in size of the cations is due to a weakening of their Lewis acidity. This is assigned to a decreasing orbitals overlap. Studies of solvent effects on basicity-dependent properties have now reached a point where a clear differentiation appears between hydrogen bond and Lewis basicities. To hammer in a nail the blacksmith and the glazier both use a hammer but not the same kind. We suggest that the diversity of the basicity phenomenon should be taken into account in future attempts to unravel solvent effects. This may become operative by the choice of the parameter (among the reliable ones) which has been defined from a property closely related to those under scrutiny. Another way to approach the various aspects of the basicity is to describe it by a multiparameter treatment. Though successfully applied to substituent effects, this matter does not yet seem mature in the field of solvent basicity. Acknowledgment. We thank Drs. R. W. Taft, M. J. Kamlet, and C. Laurence for helpful discussions and for communicating results prior to publication. We are grateful to Dr. Bollinger and to the GAF Corporation for gifts of phosphoramides and pyrrolidones, respectively. We thank Dr. J. E. Mills (Mc Neil Pharmaceutical) for providing us with a preprint on reactions of amines with methylene chloride. We gratefully acknowledge Dr. S. Geribaldi for recording the N M R spectra and Dr. M. Decouzon for purifications and analytical controls of chemicals. Registry NO. 1, 75-09-2; 2,815-24-7; 3,98-95-3; 4,75-52-5; 5,61642-2;6, 126-33-0; 7,623-81-4; 8, 100-47-0; 9,140-29-4; 10,93-58-3; 11, 78-82-0;12, 75-05-8;13, 630-18-2;14, 107-12-0;15, 766-05-2;16, 109-74-0;17, 93-89-0;18, 108-32-7;19, 616-38-6;20, 565-80-0;21, 75-07-0; 22, 107-31-3; 23, 1191-95-3; 24, 502-56-7;2k, 105-58-8; 26, 31, 123109-94-4; 27,96-22-0;28, 79-20-9; 29,75-97-8;30, 123-19-3; 36,67-64-1; 91-1;32,98-86-2; 33,563-80-4; 34, 100-52-7; 35, 141-78-6; 37,78-93-3; 38, 107-87-9; 39, 108-94-1; 40, 108-20-3; 41, 76-22-2; 42, 120-92-3;43, 502-42-1;44, 142-96-1; 45, 60-29-7; 46, 111-43-3;47, 512-56-1;48, 142-68-7;49, 592-90-5;50, 109-99-9;51, 78-59-1;52, 108-48-5;53, 80-73-9;54, 108-75-8;55, 67-68-5; 56, 2168-93-6;57, 58, 121-69-7; 59,68-12-2; 60,3772-26-7; 61,24044-24-4; 62, 632-22-4; 64,617-84-5; 65,685-91-6; 66,2591-86-8; 67, 127-19-5; 63,872-50-4; 694-85-9; 68,680-31-9; 69,4441-17-2; 70,6415-07-2; 71,109-06-8; 72, BF3, 110-86-1; 73, 108-99-6; 74, 108-89-4; 75, 121-44-8; 76,120-94-5; 7637-07-2.
