1909
THERMODYNAMICS OF ACID-BASE EQUILIBRIA
Thermodynamics of Acid-Base Equilibria.
11. Ionization of m- and
p-Hydroxybenzotrifluoride and the Concept of Fluorine
Double Bond-No Bond Resonance by Charles L. Liotta,*la D. F. Smith, Jr.,la Harry P. Hopkins, Jr.,lb and K. A. Rhodeslb School of Chemistry, Georgia Institute of Technology, Atlanta, Georgia 30338, and Department of Chemistry, Georgia State University, Atlanta, Georgia 30332 (Received January 18,1971) Publication costs assisted by the National Science Foundation and the Department of the Interior, Ofice of Water Resources Research
The free energy, enthalpy, and entropy of ionization of p- and m-hydroxybenzotrifluoride in water at 25" have been determined from pK, measurements as a function of temperature. The enthalpy of ionization was also determined by high-precision solution calorimetry. The thermodynamic data for the para and meta isomers are pK, 8.675, 8.950; AH" 4.99, 5.24 kcal/mol; and A S o -23.0, -23.3 cal/(deg mole), respectively. The results suggest that double bond-no bond resonance is unimportant in the ionization of p-hydroxybenzotrifluoride.
Recently, several workers have criticized the concept of fluorine double bond-no bond resonance as a significant influence in rate and equilibrium processes. Streitwieser aad his coworkers2 have concluded from their studies on the base-catalyzed exchange rates for replacement of hydrogen for tritium in 1H-undecaP fluorobicyclo [2.2.l]heptane and on the base-catalyzed I protodetritiation kinetics on 9-substituted 9-tritiated fluorines that fluorine hyperconjugation is not significant process have been determined in water at 25". These in stabilizing fluoroalkyl anions. From electron paraquantities have been shown to be valuable in assessing magnetic resonance studies on 2-trifluoromethylsemithe nature and degree of interaction of a substituent quinone, Stock and Suzuki3 have concluded that carwith the reaction center and the nature of the solutebon-fluorine hyperconjugation is much less significant solvent interaction. The enthalpies of ionization of than carbon-hydrogen hyperconjugation. Sheppard4 has reported that the original work by Roberts, et aZ.,j (1) (a) Georgia Institute of Technology; (b) Georgia State University. concerning the pK, of p-trifluoromethylaniliniuni ion (2) A. Streitwieser, Jr., and D. Holtz, J . Amer. Chem. SOC.,89, 692 in 50% aqueous ethanol is in error and that the up (1967); A. Streitwieser, Jr., A. P. Marchand, and A. H. Pudjaatvalue needed to place the p-CFZ substituent on a linear maka, ibid., 89, 693 (1967). free energy plot with other meta- and para-substituted (3) L. 34. Stock and J. Suzuki, ibid., 87, 3909 (1965). (4) W. A. Sheppard, ibid., 87, 2410 (1965). anilinium ions is +0.65 rather than +0.74. Stock and (5) J. D. Roberts, R. L. Webb, and E. A. McElhill, ibid., 72, 408 his coworkers6 have pointed out that a up/u,,, ratio of (1950). 1.26 f 0.01 is obtained for the trifluoromethyl deriva(6) F. W. Baker, R. C. Parish, and L. M. Stock, ibid., 89, 5677 tives of benzoic acid, aniline, and dimethylaniline, and (1967). have concluded that there is nothing unusual about the (7) Stock, et al.,e however, used the incorrect pp for the aniline syspolar contribution of the trifluoromethyl s u b s t i t ~ e n t . ~ ~tem: ~ ap/um is actually 1.33, showing that there is a small enhancement of polar effect over that predicted from the corresponding It has been reported by Jonesgthat p-hydroxybenzobenzoic acids. Stock has commented (private communication) that 1.33 is still rather close to 1.26 and that the difference does not trifluoride, when dissolved in aqueous base a t room represent a sizable enough effect to invoke the concept of fluorine temperature, is readily hydrolyzed with the loss of double bond-no bond resonance. fluoride ion.'O Roberts6 has used this observation as (8) Dewar has indicated that, in general, substituent effects are more efficiently propagated from the para position than from the evidence for charge delocalization in the anion via meta position; uP/um = 1.2: M. J. S. Dewar, "Hyperconjugation," Ronald Press, New York, N. Y., 1962, pp 159-172; M. J. S. Dewar double bond--no bond resonance interaction (I). I n and A. P. Marchand, J. A m e r . Chem. SOC.,88, 354 (1966). order to gain further insight into the extent of charge (9) R. G. Jones, ibid., 69, 2346 (1947). delocalization in this system, the thermodynamic (10) Sheppard4 has indicated that ptrifluoromethylaniline also quantities, AF", AB", and AS", for the ionization appears to decompose in aqueous solution upon standing. The Journal of Physical Chemistry, Vol. 76,No. 13,1978
1910
m- and p-hydroxybenzotrifluoride were determined by t"o methods: (a) measurement Of the p K i s a function of temperature, and (b) high-precision solution calorimetry. Experimental Section (a) pKa Measurements. The pKa of m-hydroxybenzotrifluoride in aqueous solution has been determined from 14.90 to 55.96" by a method described previous1y.l' ' Since the para isomer is known to react a t room temperature with aqueous b a ~ e , measure~!~~ ments of pH of 50% neutralized solutions were determined with respect to time using a Beckman Model 1019 pH meter and 50% neutralized p-nitrophenol as the calibration standard. I n all cases linear plots were obtained which, when extrapolated to t = 0, gave the pKa of the acid. The pKa of p-hydroxybenzotrifluoride was determined in this manner from 8.75 t o 32.80". The ionic strength of the solutions was kept low (0.0017 M ) and it was assumed that the use of 50% neutralized p-nitrophenol as the calibration standard (a compound of very similar structure to p-trifluoromethylphenol), also at the same low ionic strength (0.0017 M ) , would result in a cancellation of activity coefficients.11-13 ( b ) Calorimeter Measurements. The heats of ionization for the meta isomer were determined from the heat released when 5 ml of 4.83 N KaOH is introduced into the calorimeter containing a solution (980 ml) with a known weight of the phenol. The meta isomer was vacuum distilled on a spinning band column and stored in a sealed vial until immediately before the measurements. The observed heats for seven determinations, which have been corrected for the heat of dilution of the NaOH (y for 5 ml of 4.83 N XaOH diluted to 980 ml was calculated from "Selected Values of Chemical Thermodynamic Properties," Kational Bureau of Standards, U. S. Government Printing Office, Washington, D. C., Table 92-2, as q = 0.39 cal), are shown in Table 1. The AH of neutralization values were plotted against the square of the concentration of dissolved phenol and extrapolated to infinite dilution to obtain AR",o. The neutralization reaction carried out in the calorimeter for the para isomer involved the same type procedure except that smaller samples were used because of the appreciable heat effect accompanying the hydrolysis of the anion. The size of the sample chosen produced a heat effect well within the capabilities of the calorimeter, but did not change the drifts after the reaction by more than 35%. Within several minutes after the introduction of the sample, the thermal drifts had reached a value that was only slowly decreasing (the drifts immediately after the reaction were within 10% OS those after the second calibration some 5-8 min later). The first set of drifts that could be obtained after the reaction was used to extrapolate back to the midpoint of The Journal of Physical Chemistry, Vol, 76, N o . 13,1971
LIOTTA,SMITH,HOPKINS, AND RHODES
Table I: Heat of Neutralization D a t a for m ~ ~ y ~ r o x y ~ e n z o ~ r ~ f l uwith o r ~ ~5 eml of 4.83 N NaOH Mol X 103
door) cal
-A", kcal/mol
15.542 10.210 9.5006 4.543 4.5125 4.2189 4.376
129.8 85.1 78.4 37.3 37.1 34.9 35.9
8.39 8.33 8.25 8.20 8.23 8.26 8.22
ARN'(extrapolated) = 8 . 1 0 0.05 kcal/mol
*
the reaction which was determined from the temperature-time curve during the reaction. With this procedure, the small heat effect (increase of -0.02 cal/sec in the drifts) due to the hydrolysis was compensated for in the calculations. The corrected heats and calculated AH of neutralization value are tabulated in Table 11. The samples used in the calorimetry studies were vacuum sublimed and sealed into individual glass bulbs which were broken to introduce the sample into the solution. All of the runs were at concentrations below the lowest concentration used for the meta isomer, and the average AHN is taken to be the infinite dilution value. Table 11: Heat of Neutralization Data for p-Hydroxybenzotrifluoride with 5 ml of 4.83 iV KaOH door),
-A",
Mol X 103
cal
kcal/mol
3.132 3.774 4.125 3.606 2.962 2.479 2.607 1.897
26.89 30.42 33.88 29.99 25.26 20.48 21.30 16.47
8.58 8.06 8.21 8.31 8.53 8.26 8.17 8.68
Av = 8.35
0.06
The calorimetry values for A ~ N have ' been combined with the A B o = 13.34 kcal/mol for the ionization process of H2014,15to yield AR" for the ionization process in water of the meta and para isomers. The (11) C. L. Liotta, K. H. Leavell, and D. F. Smith, Jr., J . Phys. Chem., 71, 3091 (1967). (12) C. L. Liotta and D. F. Smith, Jr., Chem. Commun., 416 (1968). (13) W. F. O'Hara, T. Hu, and L. G. Hepler, J . Phys. Chem., 67, 1933 (1963). (14) J. Hale, R . M . Izatt, and J. J. Christensen, ibid., 67,2605 (1963). (15) C. E. Vanderzee and J. A. Smanson, ibid., 67, 2608 (1963).
THERMODYNAMICS OF ACID-BASE EQUILIBRIA
1911
AS" values were obtained from combination of AO" and A B " . The calorimeter design follows very closely that described previously.'6 A few modifications have improved the reliability and overall operation of the calorimeter. Temperature changes are followed with a set of 100-ohm thermistors placed in parallel and a highsensitivity Mueller bridge. For a temperature increase of l o ,the effective resistance of the thermistors decreases by approximately 1 ohm. I n principle, this arrangement can measure a temperature increment directly t o the nearest 0.0001". However, the heat leaks into the system were adjusted to effect a temperature change of 0.001" in a 20-30-sec time interval, making it more practical to follow the temperature to the nearest 0.001" and interpolate to the nearest 0.0001". Time intervals between successive temperatures were automatically recorded with a Sodeco printing counter activated by a switch. The calorimeter mas calibrated electrically before and after the neutralization reaction. The initial temperatures were such that the reaction midpoint was at 25 f 0.1". The most significant modification involves a unique sample introduction mechanism which eliminates the appreciable uncertainty usually associated with the determination of the heat of sample injection. I n previous calorimeters of this type, this quantity was 1-2 f 0.5 calories. The new mechanism involves a thin polyethylene bag sealed with paraffin and attached to the stirrer. The sample is introduced by turning two stainless steel knife blades a precise distance with a timing motor. The heat associated with introducing the sample with this mechanism is below the detection limits of the calorimeter, eliminating a major source of error.
Results (a) pK, Measurements. Tables III and IV summarize the pK,-temperature data for p - and m-hydroxybenzotrifluoride, respectively. I n addition, Table I11 cont,ains the change in pH with respect to time-time
Table 111: T h e 'Time and Temperature Dependence of the p H of Half-Neutralized Solutions of 0.0017 M p-Hydroxybenzotrifluoride Temp,
oc
pK, 0
8.90 10.16 12.80 17.20 20.95 22.60 28.05 29.65 32.89
8.968 8.929 8.844 8.795 8.728 8.708 8.650 8.648 8.623
----------100
200
8.963 8.920 8.836 8.780 8.700 8.673 8.571 8.561 8.494
8.958 8.910 8.827 8.764 8.670 8.636 8.495 8.472 8.364
Time, see---300 400
8.953 8.901 8.819 8.748 8.635 8.600 8.423 8.384 8.253
8.948 8.891 8.810 8.732 8.611 8.565 8.350 8.296 8.107
600
800
8.937 8.871 8.793 8.701 8.553 8.494 8.205 8.119
8.927 8.852 8.775 8.668 8.494 8.424
Table IV : The Temperature Dependence of the pK, of m-Hydroxybenzotrifluoride
15.12 15.88 16.67 17.39 18.14 18.64 19.55 20.35 21.40 22.70
9.1095 9.094 9.087 9 069 9.0615 9,049 9.0376 9.0325 9,020 8,990
23.51 24.35 25.13 26.60 27.81 28.95 30.37 31.52 32.00 32.20
I
8.982 8.960 8.952 8.939 8.914 8.898 8.885 8.876 8.871 8.867
34.62 36.82 39.12 41.40 44.50 45.63 49.50 51.20 51.60 53.07
8.833 8.804 8.784 8.757 8.734 8.712 8.687 8.669 8.667 8.654
zero being taken as the pK, at the particular temperature. Equations 1 and 2 express the pKa values of p - and pK, pK,
+ 4.949 = 1128.1/T + 5.179 =
1117.7/T
(1) (2)
m-hydroxybenzotrifluoride, respectively, as a function of absolute temperat~re.~'The pK,, enthalpy, and entropy values at 25" for each of these isomers are listed in Table V. Table V : T h e Thermodynamic D a t a a t 25' Derived from the pK,-Temperature D a t a
P K ~
P-CF~ m-CF3
8.675 8.950
AI?, kcal/mol
cal/(deg mol)
5.12 4 0 . 3 5 . 1 6 4 0.1
- 2 2 . 7 f.1 . 0 -23.7 0.4
A%",
+
( b ) Solution Calorimetry. The inherent time dependence in the pK, determinations for p-hydroxybenzotrifluoride leads to large error in the derived AZ? and AS" values, making it difficult to compare these values profitably to those of the meta isomer. The heats of ionization for the para and meta isomers were therefore investigated by means of high-precision solution calorimetry. A time dependence was also encountered in the calorimetry studies for the para isomer, but the effect on the measurement of the heat of neutralization could be minimized by the extrapolation procedures. I n this way, accurate AH" values (16) W. F. O'Hara, T. Wu, and L. G. Hepler, J . Chem. Edzlc., 38, 512 (1961). (17) Three-parameter pK,-T equations have also been derived from ~) - 58.7881 the data in Tables I and 11: (a) P K ~ ( ~ . C=F 10,511.1/T 0.108051T, (b) P K ~ ( ~ . C=F ~3051.4/T ) - 7.3937 0.020516T. (The constants were calculated by the method of orthogonal polynomial: N. W. Please, Biochem. J., 56, 196 (19541.) These authors feel, however, that the pIia-T data on these compounds are insufficiently precise to permit the use of the above nonlinear equations in 1 / T to compute the thermodynamic functions.
+
+
The Journal of Physical Chemistry. Vol. 76,No. IS, 1978
LIOTTA,SMITH, HOPKINS, AND RHODES
1912 for both isomers were obtained and combined with the ABo values obtained from the pK, measurements, yielding AS” values for both isomers to rt0.2 cal/ (deg mol). These values are tabulated in Table VI.
Table VI: The Thermodynamic Data Derived from t h e Calorimetric Measurements and the pK,’s a t 25” AI?,
P-CFB m-CFs
AS”,
PK.
koal/mol
oal/(deg mol)
8.676 8.950
4.99 i 0.08 5.24i0.08
-23.0 i 0.2 -23.3& 0.2
Discussion It has been postulated that the entropy of ionization qualitatively reflects the relative degree of solvent orientation around the un-ionized and ionized solute species involved in the equilibrium. l1,l8 For example, let us consider the ionization of p - and m-nitrophenol in water. The entropies of ionization for the para and meta isomers have been found to be - 16.9lSaand -21.1lnb cal/(mol deg), respectively. To a good approximation, it may be assumed that solvation of the un-ionized forms is the same and that entropy differences reflect the relative orientation of solvent molecules around the ionized (charged) species. While the negative charge on the oxygen of the p-nitrophenoxide ion can be delocalized over the entire molecule by direct resonance interaction with the nitro group (electromeric effect), this type of interaction should be unimportant in describing the meta isomer. As a result, one would expect the negative charge on the m-nitrophenoxide ion to be more localized on the oxygen atom. This would result in a tight solvation of the anion of the meta isomer and a comparatively loose solvation of the anion of the para isomer. This description is qualitatively reflected in the relative entropies of ionization. A comparison of the entropies of ionization of m- and p-cyanopheno120 (-21.8 and -20.0 cal/(mol deg), respectively) and m- and p-formylpheno111v21,22 (-23.4 and -20.2 cal/(mol deg), respectively) again reveals that charge delocalization of the
The Journal of Physical Chemistry, Vol. 76, No. 19,1972
para isomer produces a more positive entropy of ionieation. While Table IV reveals that the entropy of ionization of the para isomer is more positive than that of the meta isomer, the difference amounts to only 0.3 cal/ (mol deg). This small difference indicates that charge delocalization in the anion is small and that fluorine double bond-no bond resonance does not contribute significantly to this system. Analysis of the thermodynamic pK, values (2j0) determined in this study with the thermodynamic pK,’s of a wide variety of meta- and para-substituted phenols2a produces the following substituent parameters : u(*-cFs) = +0.56 and u(,+cFs) = +0.45, the ratio u ( ~ . C F ~ ) / ~ ( , - ~being F ~ ) equal to 1.25, which is essentially identical with the ratio of the corresponding benzoic acids (1.26),12a reference system in which double bondno bond resonance should be unimportant. One may conclude, therefore, that the electrical effects of the trifluoromethyl group operating in the benzoic acid system are the same as those operating in the corresponding phenolic system. I n summary, analysis of the relative entropies of ionization and relative pK, values of p - and m-hydroxybenzotrifluoride suggests that double bond-no bond resonance interaction is unimportant in the ionization of p-hydroxybenzotrifluoride.
Acknowledgment. This work was supported by the Department of the Interior, Office of Water Resources Research, as authorized under the Water Resources Research Act of 1964 and by the National Science Foundation Grant. (18) .Preprints of Papers, Symposium on Linear Free Energy Correlations, Oct 19-21, 1965, pp 94-95. (19) (a) R. A. Robinson and A. Perperl, J . Phys. Chem., 67, 2860 (1963). A value of -22.5 cal/(mol deg) has been reported by T. L. Cottrell, et al., J . Chem. Soc., 1016 (1948), but is believed to be less reliable. (b) L. P. Fernandez and L. G. Hepler, J . Amer. Chem. SOC.,81, 1783 (1959). (20) H. C. KO, W. F. O’Hara, T. Hu, and L. G. Hepler, J . Amer. Chem. Soc., 86, 1003 (1964). (21) F. J. Millero, J. C. Ahluwalia, and L. G. Hepler, J . Chem. En@. Data, 9, 319 (1964). (22) P, D. Bolton, F. M . Hall, and I. H. Reece, J . Chem. SOC.,709 (1967); Spectrochim. Acta, 22, 1825 (1966). (23) J. E. Leffler and E. Grunwald, “Rates and Equilibria of Organic Reactions,” W h y , New York, N. Y., 1963, p 373.
