NOTES
2723
T
tion of W ( E ) and the need to include the high-order Ez) if E is not very large as comterms of E,/(E pared with E,.
+
1.
Acknowledgment. The author wishes to express his sincere gratitude to Drs. A. L. Wahrhaftig and S. H. Lin for many suggestions and to Dr. C. W. Lee and Mr. D. W. Liou for discussions on computer programming. (11) E. W. Schlag, R. A. Sandsmark, and W. G. Valance, J. Phys. Chem., 69, 1431 (1965). 0.0 1.0 a.o 3.0 4.0 5 . 0 8.0 7.0 8.0 8.0 10.0
-
-'E
(ev)-+
Figure 2. W ( E ) /W,,(E) E plot for cyclopropane (E, = 2.21 ev). Frequencies 749 (2), 878 (3), 1118 (7), 1478 (3), 3221 (6) cm-1, grouped by Schlag and Sandsmark.' A, WeX&3)/W8@) calculated by Schlag and Sandsmark'.
plotted in Figure 2 and the values of the ratios of Wexact(E) as obtained by Schlag and Sandsmark? are marked with solid triangular points. From the data shown in this paper, it is apparent that the calculated values of W ( E ) by the Vestal, equation are always larger than the exact values except a t the lower energy bound ( E = 0 ev). This may be due to the approximation used in the derivation and the curve-fitted up value. The equation of Lin, et al., gave the results which are close to the exact values. It is obvious that Thiele's equation gave almost the same results as the exact values even in the low-energy region. The accuracy of the equation was also reported by Schlag, Sandsmark, and Valance." The equations of both Haarhoff and Thiele gave exactly the same results in the case of acetylene. Here, there are only four terms which need to be considered in both methods. It was found that the first four terms (v = 0, 2, 4, 6) in the two expressions are identical with each other. If a11 the rest of the terms are also equal, which they should be since the two methods involve the same basic approximation, then the two expressions are exactly identical. In the methods of Haarh ~ f f of , ~ Thiele,E and of Schlag and Sandsmark? the EZ)is important. The higher the exratio E , / ( E citation energy E the less important the high-power Ez). In Figure 2, the coincidence terms of E z / ( E of the curves of' Haarhoff and Thiele above -2 ev shows the unimportance of the high-power terms (v > 8) of E z / ( E E,) when this quantity is less than one-half. It is shown in Table I1 that the values calculated by Haarhoff's equation give much larger values of W(E)/W,,,,,(E) for t-butyl chloride (E, = 3.24 ev) than for cyclopropane (E, = 2.21 ev), but that Thiele's equation gives excellent results in both cases. This indicates the importance of E, on the calcula-
+
+
+
Infrared Spectroscopic Studies on Hydrogen Bonding between Alcohols and Ethers
by Izumi Motoyama and Charles H. Jarboel Department of Pharmacology, University of Louisville, Louisville, Kentucky (Received December 19, 1966)
Hydrogen bonding between various proton donors and ethers is well known and several ether-alcohol reactions have been studied in some detai1.2s3 Despite the volume of infrared spectroscopic work on equilibrium constants and shifts of hydroxyl group frequencies, there have been few reports dealing with the thermodynamic properties of such systems. In particular, systematic study on the thermodynamic consequences of progressive structural changes in the ether and alcohol components is lacking. We are interested in this aspect of hydrogen bonding because of its importance to drug action4 and we have examined a number of ether-alcohol systems with the objective of determining how structure affects the over-all reaction. The alcohols used were methyl, ethyl, n-propyl, isopropyl, n-butyl, isobutyl, t-butyl, and neopentyl. Ethyl and isopropyl ethers were used as the hydrogenbonding bases. This selection of substances permitted studying the synthesis of shielding and inductive effects as result from methyl substitution in both donor and acceptor. This work was carried out at concentrations prohibiting alcohol self-association and over the temperature range 21.7-48.1'. (1) To whom inquiries should be directed.
(2) (a) E. D.Becker, Spectrochim. Acta, 17, 436 (1961); (b) L. J. Bellamy, G . Eglinton, and J. F. Morman, J. Chem. SOC.,4762 (1961); ( c ) B.B. Bhowink and S. Basu, Trans. Faraday Soc., 59,813 (1963); (d) J. H.Walkup, J. Lyford, G . Marquardt, and G . W. Robinson, Trans. Kentucky Acad. Sci., 24, 101 (19G3). (3) (a) D.L. Powell and R. West, Spectrochim. Acta, 20, 983 (1964); (b) R.West, et al., J. Am. Chem. SOC.,8 6 , 3227 (1964). (4) T.Kitao and C. H. Jarboe, J. Ore. Chem., in press.
Volume 71 Number 8 July 1967 I
NOTES
2724
Table I : Hydrogen Bonding of Ethyl Ether with Alcohols
Alcohol"
Methyl Ethyl n-Propyl Isopropyl n-Butyl Isobutyl t-Butyl Neopentyl
Bonded OH frequency, cm-1 21.7O 48.1'
3493 3490 3489 3488 3492 3490 3488 3494
3500 3497 3497 3495 3498 3499 3495 3501
YAW,
21.7'
150 144 147 139 147 150 127 150
" Unbonded frequencies as rep0rted.l
em-148.1O
143 137 139 132 141 141 120 143
-e"
complex-? 21.7O 48.1'
126.9 159.4 136.0 107.8 163.0 135.1 76.6 119.6
116.7 140.0 131.4 97.1 159.4 117.8 67.9 111.7
-Kequii,
21.7O
1.27 0.79 0.91 0.83 0.73 0.89 0.74 1.03
l./mole48.1''
0.75 0.53 0.55 0.53 0.43 0.58 0.51 0.60
-AF,
-AS,
kcal/mole
eu/mole
0.14 0.16 0.06 0.11 0.18 0.07 0.18 0.02
12.2 10.4 12.3 11.o 13.4 10.7 9.5 12.9
AH,^ -kcal/mol-
3.73 2.92 3.58 3.14 3.77 3.09 2.63 3.81
(0.27) (0.15) (0.16) (0.15) (0.13) (0.16) (0.13) (0.13)
' Bracketed numbers are standard deviations. ~
Table 11: Hydrogen Bonding of Isopropyl Ether with Alcohols
Alcohol"
Methyl Ethyl n-Prop yl Isopropyl n-Butyl Isobutyl t-Butyl Neopentyl
Bonded OH frequency, cm-' 21.V 48.1°
3473 3472 3473 3476 3478 3476 3483 3484
348 1 3481 3482 3484 3484 3483 3484 3490
-Av,
21.7O
170 163 163 151 161 163 132 160
cm-148.1°
162 153 154 143 155 156 131 154
-8 complex21.7O 48.1°
126.6 136.0 140.6 75.7 171.8 147.2 42.4 119.4
110.0 119.4 135.8 68.2 167.7 127.1 30.8 113.5
-KeqUii, l./mole-
-AF,
21.7O
48.1°
kcal/mole
1.62 1.07 1.00 1.17 0.80 0.93 1.10 1.02
0.89 0.60 0.52 0.64 0.42 0.57 0.75 0.53
0.30 0.04 0.01 0.10 0.15 0.04 0.05 0.01
-AS,
-AH,*
eu/mole
-kcal/mole--
13.7 14.1 18.9 14.6 15.2 11.8 9.0 15.7
4.31 4.19 4.67 4.39 4.33 3.45 2.72 4.64
(0.14) (0.35) (0.43) (0.50) (0.18) (0.48) (0.29) (0.23)
" Unbonded frequencies as reported.& * Bracketed numbers are standard deviations.
Experimental Section All measurements were made using a Perkin-Elmer Model 137-G Infracord grating-type spectrometer. It was modified as r e p ~ r t e d ; the ~ thermostated cell systems used in these experiments were also of the type previously used. In changing cell temperature a 30min equilibration period was used. Spectra were measured a t 21.7, 30.2, 39.1, and 48.1" and average values were used when plotting Av against - A H . The solvent used in these studies was carbon tetrachloride and its desiccation was performed as r e p ~ r t e d . ~ Immediately prior to use it was passed through an 8X 2-cm column of Woelm neutral alumina in a drybox. The solute alcohols and ethers were purified by ordinary methods.6 Additionally, just prior to use the ethers were distilled from sodium in a desiccated, allglass system. Some difficulties were encountered in the purification of isopropyl ether. T o completely remove peroxides it was passed through a column of Woelm neutral alumina just before use and collected in a flask containing a small amount of the alumina. To prevent absorption of atmospheric water all The Journal of Physical Chemistry
solutions were prepared in a drybox. Solutions of the alcohols were gravimetrically prepared using samples sealed in melting-point capillaries. The capillaries were placed in desiccated 25-ml volumetric flasks containing a small volume of the appropriate ether in carbon tetrachloride. The ether-carbon tetrachloride solutions were gravimetrically prepared and were also used to dilute the alcohol samples to volume at 24". The ether-carbon tetrachloride solutions were used as the reference so that only monomeric and complexed hydroxyl bands were recorded. Alcohol concentrations were made to range between 0.002 and 0.006 M . Previous study of the alcohols has shown no self-association to occur a t this upper limit.5 The ether concentrations ranged from about 0.20 to about 0.25 M . Equilibrium constants were determined at all four temperatures and the table values taken from the limit temperatures. For each determination eight separate solutions were normally used. However, for ( 5 ) I. Motoyama and C. H. Jarboe, J . Phu8. Chem., 70, 3226 (1966). (6) J. A. Reddick and E. E. TrooDs. "Ornanic Solvents." Vol. VII.
A,' A. Weissberger, Ed., Interscience P u b h e r s , Inc., 'New York;
N. Y., 1955.
NOTES
the system isopropyl ether-methanol the number was 13, for isopropyl ether-ethanol it was 12, and for ethyl ether-methanol it was 14. At each temperature and with each solution five determinations of the spectrum below 4000 cm-l were made. To determine equilibrium constants, the alcohol monomer concentrations were calculated from the reported values of E' (apparent molar ab~orptivity).~The ea values for the hydrogenbonded complexes were determined from calculated concentrations based on the difference between alcohol added and alcohol present as monomer. Values of BO were calculated for the complexes and were shown to be proportional to ea.
Results and Discussion The data obtained are recorded in Tables I and 11. In each system studied the peaks due to alcohol monomer and complex had good symmetry a t all temperatures. Graphical resolution of the spectra demonstrated peak overlap to be insignificant; consequently, monomer concentration was calculated directly. The frequency of maximum absorbance for alcohol monomer showed no temperature dependence. The bands due to complexed alcohol showed a temperature dependence of about 7 cm-l over the range studied. The ea values for the complexes were generally much larger than those of the alcohols. They were thermosensitive and, as with ea for the alcohols, they decreased linearly as temperature advanced. The ea values for complexes of isopropyl alcohol and t-butyl alcohol with isopropyl ether were especially small. The tbutyl alcohol complexes with isopropyl ether were the only cases encountered in this study where ea values for bonded species were less than ea values of the alcohol monomer. Average values for the equilibrium constants chnracterizing formation of 1:1 alcohol-ether complexes appear in Tables I and 11. The infinitesimal decrease in free base concentration due to complexion was neglected in equilibrium constant calculations.*" It is apparent from the values a t 21.7" that isopropyl ether is a better hydrogen-bonding base than ethyl ether. The electron-releasing inductive effect of the additional methyl group adjacent to the ether oxygen is responsible for the enhanced basicity. Other than a general trend toward higher equilibrium constants when isopropyl ether is used as the base there are no readily apparent or profound structural effects on this aspect of the reaction. However, in the light of secondary steric effects it is of interest to compare the behavior of the primary alcohols. If the only structural parameters influencing equilibrium constants we1e inductive effects originating in the methyl substituents, the
2725
AKequilin going from ethyl ether to isopropyl ether should be comparable for all the primary alcohols, the change being a consequence of isopropyl ether's greater basicity. This is not the case. With ethanol the enhanced basicity of isopropyl ether is easily seen in the AKequilvalues. However, in all the other systems AKequilis small and appears to follow the degree of substitution a t the carbon atom /3 to the hydroxyl groups. We are inclined to interpret this as an incipient "parasol" effect wherein long-range conformational shielding acts to inhibit reaction with isopropyl ether. The changes in free energy, -AF, were calculated from the equilibrium constants. The enthalpy changes for complex formation, -AH, were determined from changes in the equilibrium constants at the temperatures of measurement. Entropy changes, - AS, were computed from the usual equation using the assumption that - A S is constant within the temperature range considered. The values of - A F , -AS, and - AH appear in Tables I and 11. The differences in -AH values for the two ethers is also a reflection of isopropyl ether's greater basicity. It is clear that tbutyl alcohol forms the weakest hydrogen bonds and with both ethers. The low values of -AH for reaction of this alcohol and ethyl ether is probably due to its poor proton-donating ability. The low bond strength with isopropyl ether must be due to the alcohol's low hydrogen-bonding acidity plus a direct steric effect. Molecular models bear out the presence of a large steric inhibition to bonding. By contrast, neopentyl alcohol which contains a methylene between the t-butyl group and the hydroxyl forms hydrogen bonds of ordinary strength with both ethers. A comparison of the data in Tables I and I1 shows that the average values of AV and -AH vary in the same general way. When plotted, a reasonably straight line is obtained. However, it does not fall on the Badger-Bauer' curve nor is its slope similar to that relation. It has been suggested that linear plots of AVOHand -AH for various alcohols with different hydrogen-bonding bases would show differing slopes.8 We have not found that to be true for the two ethers used in this study. Also, with the exception of tbutyl alcohol we have observed no significant decrease in AVOHas substitution about the OH group increases. This is in contrast to earlier work involving alcohols of much greater complexity.8 We relate our observation to the relative simplicity of the alcohols used.
(7) R. M. Badger and S. H. Bauer, J. Chem. Phys., 5 , 839 (1937). (8) S. Singh and C. N. R. Rao, J . Am. Chem. Soc., 88, 2142 (1966).
Volume 71, Number 8 Julv 1967
NOTES
2726
Acknowledgments. This work was supported in part by National Science Foundation Grant GB 2044, National Institutes of Health Grant GM 10363, and the Upjohn Co.
Reliability of Mole Fraction in Determining
reported by Bridger and Russell2 and Pryor, el aLa I n their system, the medium was composed of a mixture of carbon tetrachloride and a hydrocarbon. However, the composition did not change much; e g . , the largest change reported was from 27 to 55% CCl,. Table I
the Rate of Competitive Radical Reactions in Nonideal Solution
Reaction in the Absence of Diluent at 65” DMB/CsHs mole ratio
by W. J. Cheng and 31. Sswarc Department of Chemistry, State University College of Forestry at Syracuse University, Syracuse, New York 15,210 (Received January 18, 1967)
0.34 1.36 2.05 2.73 3.41
CFaH/Np
2CzFdNn
kdki
0.22 0.65 0.79 0.92 0.99
0.52 0.48 0.46 0.45 0.45
1.9 1.9 2.0 1.9 1.9
Reaction in the Presence of Diluent a t 65”
Extensive studies of competitive reactions of radicals have been carried out in this laboratory.’ These studies were concerned mainly with reactions of methyl and trifluoromethyl radicals, both in the gas phase and in solution. The investigated system basically involved two reactions described by the equations
+
+
R. HS +RH So kl (1) R . A +(AR). kz (2) In this scheme R = CH, or CFI, HS is an aliphatic
+
hydrocarbon, e.g.. 2,3-dimethylbutane (DMB) or isooctane, and A can be an olefinic, acetylenic, or aromatic compound. In deriving the k2/kl values from the experimental data it is assumed that the rates of reactions 1 and 2 are proportional to the mole fractions of HS and A, respectively. This assumption is fully justified for a gas phase reaction; however, one could question its validity when the process takes place in the liquid phase, particularly if the relevant components do not form an ideal solution. Several methods are available to test this assumption. For example, one may vary the composition of the mixture from that rich in HS to that rich in A and investigate the constancy of the calculated k2/k1 values. In most of our studies this could not be achieved since the usually high reactivity of the selected substrates forced us to-work at high-dilutions of A. The present work a system in which the composition of the reacting substrates, 2,3-dimethylbutane ( D ~ I B )and benzene, could be varied from a DMB/C& mole ratio of 0.34 to 3.4. As shown by the data in Table I, this tenfold variation in the mole ratio did not affectthe respective k2/kl value. Similar conclusions may be deduced from the results The Journal of Physical Chemistry
+
(DMB c~H~)~/c~F~Q volume ratio CFaH/Ns
1:9 1:9 5:5 5:5 9:1 9:1
0.80 0.795 0.81 0.825 0.81 0.81
2CzFe/Nz
0.48 0.47 0.41 0.38(?) 0.38(?) 0.44
kdki’
1.8 1.9 2.0 2.0 2.0 1.9
DMB/C& mole ratio was kept constant a t 2.05. * C6F12 = dimer of perfluoropropylene involving a cyclobutane ring. ’ kz/kl = { [2 - (CFaH/Nz) (2CzF,/Nz)I/(CF,H/Nz)J(DMB/ CsHe).
-
It should be stressed that in such a study the extent of the cage reaction has to be determined directly for each composition, because its probability varies with the composition of the liquid (see the column 2CzF.y‘Nz in Table I), The variation in the extent of the cage reaction affects the calculated amount of “CFaH lost” and this latter value is required for the calculation of kz/ki. A more severe test of the validity of the above assumption involves experiments in which an inert liquid is added to the mixture of the reagents. This was done in the second series of runs reported in Table I. The mole ratio DMB/CeHs was kept constant, (1) (a) M. Szwarc and J. H. Binks, “Theoretical Organic Chemistry,” Kekule Symposium, London, 1958, Butterworth and Co. Ltd., 1959, pp 147-186; (b) A. P. Stefani and M. Szwarc, Proceedings of the Second International Symposium on Fluorine Chemistry, Colorado University Press, Boulder, Colo., 1962, pp 334-323; (c) M. Szwarc, Special Publication No. 16, The Chemical Society, London, 1962, p 91. (2) R. F. Bridger and G. A. Russell, J. A m . Chem. SOC.,85, 3754 (1963). (3) w.A. Pryor, J. T.Echois, and K. Smith, ibid., 88, 1189 (1966).
