J. Phys. %hem. 1985, 89. 3347-3352 at the highest concentration of DMPO used (1.5 X M in our experiment) is not sufficiently fast enough for the deconvolution procedure to effect a significant difference between the T ~ ' Sand q ; s when the response time of the detection system is not as yet approached. Figure 6 is a plot of ( ~ ~ ) -vs. l [DMPO]. The slope of the straight line through the origin gives the bimolecular trapping rate constant k3 as (1.1 f 0.1) X lo8 M-' s-'. This is in excellent agreement with the results obtained at 100-kHz field modulation. DMPO has been shown to be more efficient than the,acyclic spin trap phenyl-N-tert-butylnitrone(PBN). Thus InggldEand co-workers have determined the absolute trapping rate constants of primary and secondary alkyl radicals by DMPO in benzene at 40 OC to be 26 X lo5and 4.2 X los M-' s-', respectively. For PBN, the corresponding absolute trapping rate constants are 1.3 X lo5 and 0.68 X los M-' s-l, respectively. Greenstock and Wiebels have determined by pulse radiolysis the trapping rate constant of acetone ketyl radical by PBN to be 1.0 X lo7 M-' s-'. Our measured value of 1.1 X lo8 M-' s-' for trabping of acetone ketyl radicals by DMPO is 10 times larger than the corresponding rate constant for trapping of acetone ketyl radicals by PBN. This exhibits the same trend as that found for the previously mentioned reactions studied by Ingold.E
-
3347
Conclusions In this time-resolved EPR study of a spin-trapping reaction, the deconvolution procedure seems able to recover reasonable rise times from highly distorted kinetic profiles. The results obtained at 100-kHz field modulation compare favorably with those a t 2-MHz field modulation. With a careful choice of the form of R(t) one can measure reliably kinetic rate processes with rise times which are comparable to the instrument response time of the EPR spectrometer. Although we have developed and applied this method for FPEPR kinetic studies, the general method should be applicable to the kinetic response of any detection system.17
Acknowledgment. This work was supported financially by an Operating Grant from the Natural Sciences and Engineering Research Council of Canada. Registry NO. CH$(=O)CH,, 67-64-1; CH,CH(OH)CH3,67-63-0; DMPO,3317-61-1; SA, 74670-74-9; CH,C(OH)CH3, 5131-95-3. (17) A referee has pointed out that a similar treatment to correct for the effects of the pulse profile and instrumental characteristics has been applied to pulse radiolysis kinetic profiles.'* (18) Janata, E.; Schuler, R. H. J. Phys. Chem. 1982,86,2078. Scaiano, J. C. J . Phofochem. 1981, 16, 71.
Gas-Phase 'H NMR Studies of Keto-Enol Tautomerism of Acetylacetone, Methyl Acetoacetate, and Ethyl Acetoacetate Michael M. Fokendt, Boris E. Weiss-Lopez, J. Paul Cbauvel, Jr., and Nancy S. True* Department of Chemistry, University of California, Davis, California 9561 6 (Received: November 27, 1984)
The effect of intermolecular interactions on the keto-enol tautomeric equilibria of acetylacetone, ethyl acetoacetate, and methyl acetoacetate has been studied by using 'HNMR spectroscopy. For all three molecules gas-phase spectra obtained in the 445-373 K temperature range are consistent with equilibrium constants which favor the enol tautomer to a larger extent than is found in condensed phases at corresponding temperatures. Temperature-dependent intensity ratios yield the following gas-phase relative enthalpies (W(eno1-keto), kcal/mol): -4.66 (18), -3.17 (24), and -3.04 (59) for acetylacetone, ethyl acetoacetate, and methyl acetoacetate, respectively. The corresponding differences in neat liquid samples are 2-3 kcal/mol smaller. For all three molecules in all phases the keto tautomer is favored entropically (ASo(enol-keto) is -8 cal/(mol K)). These results are discussed in terms of structural differences and dielectric effects.
Introduction For gaseous molecules, N M R spectroscopy can provide high quality thermodynamic parameters characterizing equilibria in suitably selected cases.' Comparison with similarly obtained condensed-phase data yields a straightforward measure of the direction and extent of solvent effects on both relative enthalpies and relative entropies of the associated molecular forms. The present study reports gas-phase thermodynamic data for the keto-enol equilibria of acetylacetone (2,4-pentanedione) (R, =
R2 = CH,), methyl acetoacetate (R, = CH3, R2 = OCH,), and ethyl acetoacetate (R, = CH3, R2 = OC,H,). The high energy bamer to tautomer interconversion permits observation of distinct N M R spectra of both forms at temperatures well above ambient thus allowing quantitative gas-phase thermodynamic data to be obtained for the first time for these relatively nonlabile molecules. In each case a comparison with condensed-phase data reveals significant phase-dependent relative energy and relative entropy (1) Chauvel, J. P., Jr.; True, N. S. J . Phys. Chem. 1983,87, 1622-1625.
0022-3654/85/2089-3347$01.50/0
differences between the keto and enol tautomers. The keto-enol equilibrium of acetylacetone has been well characterized in solutions by using 'H N M R spectroscopy. The relative thermodynamic parameters are found to be both solvent and concentration dependent. Recent temperature-dependent studies of acetylacetone in seven solvents in the range 273-353 K demonstrated that in all cases the enol form is favored energetically (AHo298(E-K) ranged from -1.8 (Me2S0 solution) to -3 kcal/mol (cyclohexane solution)) and disfavored entropically (hS0298(E-K) ranged from -4 (cyclohexane solution) to -6 cal/(mol K) (neat liquid)).2 These results are consistent with several earlier N M R studies of equilibria in a~etylacetone.~-~ In contrast to the liquid-phase system, only semiquantitative data concerning the gas-phase keto-enol equilibrium of acetylacetone exists. Spectroscopic&' and electron d i f f r a c t i ~ n ' ~studies J ~ have, (2) Spencer, J. N.; Holmboe, E. E.; Kirshenbaum, M. R.; Firth, D. W.; Pinto, P. B. Can. J . Chem. 1982, 60, 1178-1182. (3) Reeves, L. W. Can. J . Chem. 1957, 35, 1351-1365. (4) Burdett, J. L.; Rogers, M. T. J . Am. Chem. SOC.1964, 86, 2105. (5) Rogers, M. T.; Burdett, J. L. Can. J . Chem. 1965, 43, 1516-1526. (6) Powling, J.; Bernstein, H. J. J. Am. Chem. Soc. 1951, 73,4353-4356. (7) Jarrett, H. S.; Sadler, M. S.;Shoolery, J. N. J. Chem. Phys. 1953.21, 2092-2093. (8) Harris, R. K.; Rao, R. C. Org. Magn. Reson. 1983, 21, 580-586. (9) Funck, E.; Mecke, R. In "Hydrogen Bonding"; Hadzi, D., Ed.; Pergamon Press: London, 1959; p 433.
0 1985 American Chemical Society
3348 The Journal of Physical Chemistry, Vol. 89, No. 15, 1985 however, established that the enol form of acetylacetone also predominates in the gas phase. An equilibrium constant (Kq = enol-keto) of 2.6 was obtained a t 419 K from previous gas-phase N M R measurements,* of 9 at 373 K from gas I R spectral meas u r e m e n t ~and , ~ of l .9 (0.3) at 378 K from electron diffraction measurements.12 An earlier electron diffraction study at 290 K was consistent with a K of 32.13 A UV study a t 298 K did not detect absorptions from &e keto form.’O The observed ionization energies from X-ray photoelectron spectra of acetylacetone a t temperatures between 323 and 423 K were consistent with expectations for the enol form and were not temperature dependent, indicating the presence of a single major tautomer.” All these data demonstrate that the enol form is predominant in the gas phase, but they are not of sufficient quality to yield quantitative values for the thermodynamic parameters characterizing the equilibrium. Substitution of an 0-alkyl group for a methyl group dramatically reduces the Kq value for the tautomer equilibria. At 306 K liquid methyl acetoacetate was found to exist exclusively in the keto form while ethyl acetoacetate produced ‘H N M R spectra consistent with a K,(enol/keto) of 0.64 (0.1 mole fraction in hexane), and 0.081 for the neat l i q ~ i d . ~ The present study reports thermodynamic parameters for keto-enol equilibria in three molecules and trends in their associated gas-liquid differences. Factors contributing to the observed entropy and enthalpy differences in the gas phase are discussed as well as expected solvent perturbations on these systems based on solvent packing and electrostatic forces.
Experimental Section The following chemicals were purchased from the sources indicated and were used without additional purification: acetylacetone (Aldrich, gold label), ethyl acetoacetate (Aldrich), methyl acetoacetate (Sigma, Grade 1, 99%+ purity), and CHCl, (Mallinkrodt, AR grade). Gaseous chloroform was used as a frequency and resolution standard. Its relative chemical shift, 7.102 ppm downfield from gaseous Me4Si, is pressure and temperature independent a t pressures between 1 and 5000 torr and temperatures between 240 itnd 450 K. Gas-phase N M R sample mils were constructed from Wilmad 12-mm N M R tubes as previously d e ~ c r i b e d . ’ ~Approximately 6 pL of the liquid dione and 2 pL of chloroform were injected into the N M R cells and were subsequently thoroughly degassed and sealed with a torch to yield sample partial pressures of -200 torr of dione and 100 torr of chloroform upon complete vaporization at -425 K. Neat liquid samples were made in 8-mm-i.d. 12-mm-0.d. high-pressure N M R tubes (Wilmad, special order), Neat liquid samples were thoroughly degassed and sealed with a torch. Liquid chloroform was used as a resolution and frequency reference. The chemical shift of chloroform in solution is solvent, concentration, and temperature dependent, with a value of 7.258 ppm downfield of Me4Si being used for the reference chemical shift. The solvent used in previous studies, carbon tetrachloride, could not be used over the temperature range studied in this work due to its low boiling point (349.9 K). Naphthalene is a suitable solvent for nonpolar studies (boiling point 491.1 K). A 1% acetylacetone in naphthalene solution was used to obtain thermodynamic data. All NMR measurements were made with a Nicolet NT 200 FT spectrometer with proton observation at 200.067 MHz using a 12-mm IH probe. All measurements were made on nonspinning samples in the unlocked mode. For the gas-phase N M R samples the restricted volume N M R tubes were placed in the probe by using a 20-mm spinner fitted with a Teflon adapter. The restricted
-
(10) Cheng, L. T., Ph.D. Dissertation, Louisiana State University, Baton Rouge, LA, 1968. (11) Houk, K. N.; Davis, L. P.; Newkomc, G. R.; Duke, Jr., R. E.;Nauman, R. V. J . Am. Chem. Soc. 1973, 95, 8365-8371. (12) Lowery, A. H.; George, C.; D’Antonia. P.; Karle, J. J . Am. Chem. SOC.1971, 93,-6399-6403. (13) Andreassen, A. L.; Bauer, S. H. J. Mol. Srrucr. 1972, 12, 381-393. (14) Chauvel, J. P., Jr.; Friedman, B. Ri.; Van, H.; Winegar, E. D.; True, N. S.J . Chem. Phys., in press.
Folkendt et al. volume tubes were necessary to eliminate convection and minimize temperature gradients in the cell. Gas-phase proton spectra were obtained between 372.8 and 444.9 K. The temperature was calibrated by using an empty (1 atm air) cell containing four copper constantan thermocouples. Typical gradients range from 0.1 K at 373 K to 0.3 K at 445 K across the 12-mm portion of the tube. Liquid-phase spectra were taken by inserting the 12mm-0.d. 8-mm-i.d. N M R tubes in the 12-mm spinner. Liquidphase proton spectra were obtained between 328 and 423 K. The temperature was calibrated by placing four copper constantan thermocouples in the active volume of a N M R tube filled with ethyl acetoacetate. Typical temperature gradients across the active volume ranged from 0.2 K at 328 K to 0.5 K at 423 K. Before acquiring transients, each sample was allowed 15 min to equilibrate. Measurements of the keto-enol equilibrium constant of acetylacetone at 373 K every 30 min over a period of 3 h showed no change within its error limits. Similar measurements were made with ethyl acetoacetate. Gas-phase spin-lattice relaxation times for acetylacetone and ethyl acetoacetate were determined by the inversion recovery method. Typical Tl’s ranged from 1.5 to 2.1 s for the methyl, metliylene, and vinyl protons. The hydroxylic protons had Tl’s of 54 ms in acetylacetone and 457 ms in ethyl acetoacetate. Gas-phase proton spectra were taken by accumulating -60 transients into 16K of memory. A 67’ pulse with a 3600-Hz sweep width was used. The delay between pulses was 16 s, which is greater than 5Tl’s of the longest TI of each molecule. The transients were subjected to sensitivity enhancement (line broadening 3 Hz) and Fourier transformed into the frequency domain with signal td noise ratios greater than 500/1. The spectra were referenced to gaseous chloroform (7.102 ppm) which in turn was referenced to gaseous Me4Si (0.00 ppm). Liquid-phase ‘H N M R spectra were acquired as described above, except that a low power pulse corresponding to a 2O flip angle was used. Typically, 16 transients were summed, multiplied by a 1-Hz sensitivity enhancement, and Fourier transformed to produce frequency domain spectra with signal to noise ratios greater than 500/1. The liquid-phase spectra were referenced to chloroform, which is -7.258 ppm downfield of Me.&. Previous liquid-phase spectral assignments for acetylacetone,”’ methyl acet~acetate,~ and ethyl acetoacetate5 were found to apply to gas-phase spectra. For each molecule, the vinyl resonance was integrated to determine the population of the enol form and the methylene resonance was integrated to determine the population of the keto form. The methylene integrals were divided by a factor of 2 prior to taking the ratio Kq = enol/keto in order to normalize for the number of protons responsible for each resonance. The relative intensity data was obtained via -7 independent integrations a t each temperature. The mean value of the keto-enol equilibrium constant, Kq = enol/keto, was determihed at 19 temperatures by this method. The resulting temperature dependence of Kq was analyzed via a linear regression fit to the van’t Hoff equation by using the program MINI TAB'^ to obtain thermodynamic parameters. Weights were assigned in the usual manner.I6 All values reported in Table I are at the 95.5% confidence limit.
-
Results Table I summarizes thermodynamic parameters characterizing the ketc-enol equilibrium of acetylacetone, methyl acetoacetate, and ethyl acetoacetate. Supplementary Tables 1-3 contain complete equilibrium constant data for acetylacetone, ethyl acetoacetate, and methyl acetoacetate, respectively. Temperature-dependent ‘H NMR spectra, thermodynamic results, and spin relaxation data (TI and T2) are summarized below. Acetylacetone. Figure 1 displays representative ‘H NMR spectra of acetylacetone gas (a), 1% solution in naphthalene (b), (15) Minitab Project, Statistics Department, 21 5 Pond Laboratory, PennsylvaniaState University. (1 6) Bevington, P. R. “Data Reduction and Error Analysis for the Physical Sciences”; McGraw-Hill: New York, 1969.
The Journal of Physical Chemistry, Vol. 89, No. 15, 1985 3349
Keto-Enol Equilibria
Figure 1. Temperature-dependent ‘H NMR spectra for acetylacetone: (a) 200 torr of gas, (b) 1% liquid in naphthalene, (c) neat liquid. The labels
-
refer to temperatures in K. The resonances at -5.3, 3.3, 2.0, and 1.9 ppm are due to the enol vinyl proton, the keto methylene protons, and the enol and keto methyl protons, respectively. The enol hydroxylic proton resonance, which is at 15 ppm, is not shown.
TABLE I: Cas- and Liquid-Phase Thermodynamic Parameters (Enol-Keto) for @-DicarbonylCompounds Derived from a van’t Hoff Fit to Integrated ‘H N M R Spectra” AG0298(E-K)
AZfo(E-K)
ASo(E-K)
Acetylacetone gas neat liquid
-2.20 (45)
-4.66 (18)
-8.26 (45)
b
-0.91 (30)
-2.81 (12) -2.8
-6.39 (30) -6.4
-1.14 (58)
-2.47 (23)
-4.46 (60)
C
1% solution in naphthalene
Ethyl Acetoacetate gas neat liquid b
d
-0.08 (59)
-3.17 (24)
-10.38 (57)
1.55 (51) 1.51
-0.18 (20)
-5.81 (54)
b
L1D
La
LD
zm
zm
zm
lDDDn
Methyl Acetoacetate gas neat liquid
LB
-0.08 (1.45)
-3.04 (59)
-9.93 (1.43)
1.89 (1.61)
0.08 (62)
-6.07 (1.72)
“AG029s(E-K) and M0(E-K) are in units of kcal/mol, while ASo(E-K) is in cal/(mol K). All uncertainties are reported to 2a. bSolutions study, this work. CSolutionstudy, ref 4. dSolution study, ref 2.
and neat liquid (c) between 377.0 and 417.6 K. In the gas phase, the resonances at 5.289 and 1.867 ppm are due to the enol form vinyl and methyl group protons, while those at 3.299 and 2.017 ppm are due to the keto form methylene and methyl group protons. In the liquid phase the chemical shifts are temperature, solvent, and concentration dependent. However, in each case the resonances at ca. 5.17, 3.21, 1.78, and 1.63 ppm are assigned to the enol vinyl proton, the keto methylene protons, the enol methyl protons, and the keto methyl protons, respectively. The hydroxylic proton resonances are 15.2, 14.6, and 15.7 ppm downfield from TMS, in the gas, neat liquid, and 1% solution in naphthalene, respectively. In the gas phase the hydroxyl proton rwnance fwhm is -16.7 Hz a t 373 K, -28 Hz at 423 K, and -41 Hz a t 445 K, significantly broader than the line widths of the other proton resonances of acetylacetone (-2 Hz). The spectra shown in Figure 1 demonstrate that the enol form predominates in all phases and at all temperatures where intensity measurements were obtained. Also, its relative population is inversely proportional to temperature, demonstrating that it is more stable than the keto form. The temperature dependence is greater in the gas phase. Nineteen temperaturedependent intensity ratios
Figure 2. van’t Hoff plot for acetylacetone: (A) 200 torr of gas, (B) 1% in naphthalene, (C) neat liquid. The associated thermodynamic parameters are in Table I. Equilibrium constant data are in supplementary Table 1. for the gas obtained between 372.8 and 444.9 K, 19 temperature-dependent intensity ratios obtained between 366.3 and 423.2 K for the neat liquid, and 18 temperature-dependent intensity ratios obtained between 366.3 and 423.2 K for the naphthalene solution are shown graphically in the van’t Hoff plots shown in Figure 2. Equilibrium constants appear in supplementary Table I. The gas-phase experimental temperature range was limited at the low end by volatility and a t the high end by N M R probe specifications. Liquid-phase temperature ranges were limited at the low end by the length of time required for the samples to reach thermodynamic equilibrium (- 15 min at 373 K, becoming greater than 2 h a t 300 K), and a t the high end by the boiling points of the various solutions. The results appearing in Table I demonstrate that the enol form of acetylacetone is energetically favored in all phases and has significantly less entropy than the keto form. Both effects are more pronounced in the gas phase. The entropy of the two tautomers is more similar in the naphthalene solution. The values of T1for 200 torr of acetylacetone at 423 K are 1.7 1, 1.53, and 1.1 3 s for the enol vinyl resonance, the keto methylene resonance, and the methyl resonances, respectively. The value of T I for the enol hydroxylic resonance is only 54 ms a t 423 K. Ethyl Acetoacetate. Figure 3 displays representative ‘H N M R spectra of 100 torr of ethyl acetoacetate gas. The resonances at 12.01,4.846, and 1.814 ppm are due to the enol hydroxylic proton (not shown), vinyl proton, and methyl protons (not shown), re-
3350
The Journal of Physical Chemistry, Vol. 89, No. 15, 1985
Folkendt et al.
t n i U , L
-2.m
*
ZB
*
- -_- -
'
j
2.
4
3
.
.
'
*
I
r I
zm
z. n
u law
402 0
5
*
4
I) 8
Figure 4. van't Hoff plot for ethyl acetoacetate: (A) 100 torr of gas, (B) neat liquid. The associated thermodynamic parameters appear in Table I. Equilibrium constant data are in supplementary Table 2.
PPM
Figure 3. Temperature-dependent 'H N M R spectra of 100 torr of gaseous ethyl acetoacetate. The labels refer to temperatures in K. The resonance at -3.5 ppm is due to a volatile impurity. The resonances at -4.8, 4.2, and 3.1 ppm are due to the enol vinyl proton, the ester side chain methylene protons, and the keto methylene protons, respectively. The enol hydroxylic resonance, which is at -12 ppm, and the methyl resonances are not shown. spectively. The resonances a t 3.147 and 2.063 ppm are due to the keto methylene and methyl protons (not shown), respectively. The ethyl ester moiety (CH2, 4.153 ppm; CH3, 1.218 ppm (not shown)) of both tautomers is magnetically indistinguishable. A volatile impurity in the gas phase with a relative concentration of -5% has two resonances, one at 3.633 ppm and one at 1.917 ppm (not shown). In the neat liquid at 423 K ethyl acetoacetate has resonances at ca. 11.6, 4.61, 3.80, 3.03, 1.82, 1.56, and 0.88 ppm which are due to the enol hydroxylic proton, the enol vinyl proton, the ester methylene protons, the keto methylene protons, the enol methyl group, the keto methyl group, and the ester methyl group, respectively. These resonances are temperature, solvent, and concentration dependent. The volatile impurity which was present in the gas phase is absent in the-liquid phase spectra. The values of T I for ethyl acetoacetate have been determined for a 100-torr sample at 423 K. For the keto form they are 1.88 and 1.84 s for the methylene and methyl protons, respectively. The ester side chain has Tl's of 2.13 and 1.61 s for the methylene and methyl protons, respectively. The enol form has T1'sof 2.20 s for the vinyl resonance, and 0.457 s for the hydroxylic resonance. It should be noted that the value of Tl for the hydroxylic proton in ethyl acetoacetate is 10 times that of the acetylacetone at similar temperatures and pressures. The natural line width for this resonance in ethyl acetoacetate is only 4.4 Hz, indicating that the T I contributions to T2 are small. The temperature range studied was limited by sample volatility, the amount of time required to achieve thermodynamic equilibrium, and NMR probe design. K, values appear in supplementary Table 2. Figure 4 shows a van't Hoff plot for the gas and neat liquid. The associated thermodynamic parameters are in Table I. The data demonstrate that the enol form of ethyl acetoacetate is more stable in the gas phase and is slightly more stable in the liquid phase. The keto form is favored entropically in both phases. Merhyl Acetoacetate. Figure 5 displays representative 'H NMR spectra for 100 torr of gaseous methyl acetoacetate at -405 K. In the gas phase a t 423 K the resonances at 11.97,4.865, and 1.819 ppm are due to the enol hydroxylic proton (not shown), vinyl proton, and methyl protons (not shown), respectively. The resonances at 3.155 and 2.059 ppm are due to the keto methylene and methyl protons (not shown), respectively. The methyl ester moiety of both tautomers is magnetically equivalent ('H methyl
-
5
4
3
PPM
Figure 5. Temperature-dependent 'H N M R spectra for 100 torr of gaseous methyl acetoacetate. The labels refer to temperatures in K. The resonance at -3.4 ppm is due to a volatile impurity, while the resonances at -4.8, 3.6, and 3.2 ppm are due to the enol vinyl proton, the ester side chain methyl protons, and the keto methylene protons, respectively. The enol hydroxylic proton resonance, which is at 12 ppm, and the enol and keto methyl resonances are not shown. The plot has been scaled so that the keto methylene resonance is full scale.
-
-
resonance, 3.641 ppm). Once again, there is a volatile impurity which has a relative concentration of 10% in the gas-phase sample, with resonances at 3.410 and 1.920 ppm (not shown). Proton NMR spectra of neat liquid methyl acetoacetate at 423 K show monances at 11.96,4.98, 3.77, 3.42, and 2.17 ppm which are assigned to the enol hydroxylic proton, the enol vinyl proton, the methyl ester group protons, the keto methylene protons, and the keto methyl group protons, respectively. The chemical shifts for neat liquid methyl acetoacetate are temperature, solvent and concentration dependent. The values of K,, obtained from the 'H spectra in Figure 5 , are in supplementary Table 3 and are plotted for the gas and neat liquid in a van't Hoff plot in Figure 6. The resulting thermodynamic parameters for methyl acetoacetate are in Table I. The enol tautomer is favored enthalpically, while the keto form is favored entropically. The values of AHo for ethyl acetoacetate and methyl acetoacetate are identical within error limits.
Discussion The significant results of this study include the greater stability of the enol tautomers and the greater entropy of the keto tautomers
The Journal of Physical Chemistry, Vol. 89, No. 15, 1985 3351
Keto-Enol Equilibria
1 -..
-
-210
-
-2.
-
i Z.23
La
am
2.6
IDM
Figure 6. van’t Hoff plot for methyl acetoacetate: (A) 100 torr of gas, (B) neat liquid. The associated thermodynamic parameters appear in Table I. Equilibrium constant data are in supplementary Table 3.
for all three 0-diketones, the similarity of the thermodynamic parameters for both 8-diketoesters, and the similar solvent effects on the thermodynamic parameters. These topics are discussed below. Interesting qualitative differences in spin-lattice ( T I ) relaxation rate constants of hydroxylic protons of symmetrically and asymmetrically substituted enols are also discussed. Table I indicates that in the gas phase the enol tautomers of acetylacetone, ethyl acetoacetate, and methyl acetoacetate are favored enthalpically over the keto tautomers. Theoretical calculations of relative stabilities of keto-enol tautomers are not extensive. Most theoretical studies have concerned the nature of the intramolecular hydrogen bond in the enol form of malonaldehyde (1,3-propanedi0ne).~’-~~An early M I N D 0 calculation for acetylacetone places the energy difference between the enol and keto tautomers at 0.47 kcallm01.~~Structural data available for both tautomers of acetylacetone and experimental and theoretical studies of the vibrational dynamics of the intramolecularly hydrogen-bonded enol tautomer of acetylacetone and malonaldehyde do, however, provide insight into the significantly greater stability of the enol tautomer. Electron diffraction studies of acetylacetone have demonstrated that the enol tautomer has a planar ring structure and that the keto form consists of two planar acetyl groups with a dihedral angle T ( ~ of ~49 ~(4)O.I2 ) Microwave spectra of isotopically substituted forms of malona l d e h ~ d e ~ are ~ - ~also ’ consistent with a planar C, ring structure containing an intramolecular hydrogen bond. The inclusion of an intramolecular hydrogen bond and concurrent resonance stabilization lowers the internal energy of the enol form. The results which appear in Table I also demonstrate that the keto tautomers of @-decarbonylcompounds have significantly more entropy than the enol tautomers. It is straightforward exercise in statistical mechanics to estimate the entropy difference between the two tautomers assuming reasonable models for the large-amplitude internal motions unique to each form. The entropy differences depend on the difference between the entropy associated (17) Bicerano, J.; Schaefer 111, H. F.; Miller, W. H. J . Am. Chem. SOC. 1983, 105, 2550-2553. (18) Kato, S.;Kato, H.; Fukui, K. J . Am. Chem. Soc. 1977,99,684-691. (19) Karlstrom, G.; Wennerstrom, H.; Jonsson, B.; Forsen, S.; Almlof, J.; Roos, B. J . Am. Chem. SOC.1975, 97, 4188-4192. (20) Karlstrom, G.; Jonsson, B.; Roos, B.; Wennerstrom, H. J. Am. Chem. SOC. 1976, 98, 6851-6854. (21) Bouma, W. J.; Vincent, M. A.; Radom, L. Int. J . Quantum Chem. 1978, 14, 761-711. (22) Del Bene, J. E.; Kochenour, W. L. J. Am. Chem. SOC.1976, 98, 204 1-2046. (23) Dewar, M. J. S.; Shanshal, M. J . Chem. SOC.A 1971, 25-29. (24) Rowe, W. F., Jr.; Dutrst, R. W.; Wilson, E. B. J . Am. Chem. SOC. 1976, 98, 421-4023, (25) Baughcum, S. L.; Duerst, R. W.; Rowe, W. F.; Smith, Z.; Wilson, E. 8. J . Am. Chem. Soc. 1981, 103, 6296-6303. (26) Baughcum, S. L.; Smith, Z.; Wilson, E. B.; Duerst, R. W. J . Am. Chem. SOC.1984, 106, 2260-2265. (27) Turner, P.; Baughcum, S. L.; Coy, S.L.; Smith, Z. J. Am. Chem. Soc. 1985, 106, 2265-2261.
with the low-frequency torsions of the keto form and the entropy associated with the large-amplitude internal H motion and out of plane torsions of the ring in the enol form, assuming that all other vibrational frequencies are the same in both tautomers. Estimates of the entropy contributions from these internal motions are described in the following two paragraphs. Since vibrational frequencies for the torsional modes and for the hydrogen internal motion have not been observed experimentally, these estimates are quite qualitative. Both and experimental2e27 studies demonstrate the formation of an intramolecular hydrogen bond in the enol form. The associated potential function has been studied theoretically and experimentally for malonaldehyde and experimentally for acetylacetone with a general conclusion that a C, structure is a minimum. The most recent theoretical results for malonaldehyde at a slightly better than double basis” yield a barrier height of 8 kcal/mol, -2 kcal/mol lower than that from earlier theoretical Very recently the tunnelling splitting, 21.583 cm-l, has been determined from microwave spectroscopic measurements, allowing limits on the barrier height (4-7 kcal/mol) and the distance between the minima (0.994-0.822 A) to be ~ b t a i n e d . ~ ~ ? ~ ’ These results are consistent with energy levels for the next two tunnelling states of between 860 and 1420 cm-I and for the second pair between 1190 and 1820 cm-’. The entropy associated with this vibration accordingly is small, ranging from - 1 to 0.3 cal/(mol K) a t 400 K. Liquid-phase NMR studies of acetylacetone which determined the quadrupole coupling constant for the bridging hydrogenZBare also consistent with a double minimum potential function for the intramolecular hydrogen bond. Significant differences in the details of the potential function of malonaldehyde, acetylacetone, and other substituted P-dicarbonyls are unlikely. The two lowest out of plane vibrations of malonaldehyde are calculated from derivatives of the potential as 261 and 255 cm-’ and are consistent with existing vibrational data.29 Associated entropies for each torsion evaluated at 400 K are -2.5 cal/(mol K). Less information is available concerning the torsional potentials of the keto tautomer. Electron diffraction results are consistent with a nonplanar structure having the acetyl groups rotated 48.6 (4)O with respect to each other. It cannot be concluded from this information if this thermal average structural arrangement does in fact correspond to a minimum in the internal rotation potential function. Based on previous results for similar compounds, which are described below, it can be reasonably assumed that the torsional barriers are -2.5-5 kcal/mol, consistent with frequencies which are considerably lower than the enol out of plane vibrations. Propionaldehyde has an internal rotation barrier about the carbonyl carbon-methylene carbon bond of 2 k c a l / m 0 1 ; ~ pro~*~~ pionyl fluoride32 has an analogous internal rotation potential function. For both molecules, the most stable conformer has the alkyl moiety oriented syn to the carbonyl bond with skew (torsional angles of 120O) conformers present at higher ( 1 kcal/mol) relative energies. Since acetylacetone has torsions about similar bonds (adjacent to COX moieties) but with bulkier substituents, we estimate that its torsional barriers are slightly larger than those in propynal and propionyl fluoride. In order to calculate the partition function associated with internal rotation in this system the potential was expanded in a Fourier series and the Schrodinger equation was solved by using the program of Laane et al.33 The calculation was performed for different values of the high barrier (2.5-5 kcal/mol) and for different shapes of the potential function. The entropy contribution for each torsion ranges from 6.5 to 7.3 cal/(mol K) and is relatively insensitive to the actual form of the potential. It is clear, however, that low-frequency torsional vi-
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(28) Egan, W.; Gunnarsson,G.; Bull, T.E.; Forsen, S. J . Am. Chem. SOC. 1977, 99,4568-4572. (29) Ogoshi, H.; Nakamoto, K. J . Chem. Phys. 1966, 45, 3113-3120. (30) Butcher, S. H.; Wilson, E. B. J . Chem. Phys. 196440,1671-1678. (31) Pickett, H. M.; Scroggin, D. G. J . Chem. Phys. 1974,61,3954-3958. (32) Stiefvater, 0. L.; Wilson, E. B. J. Chem. Phys. 1969,50, 5385-5403. (33) Lewis, J. D.; Malloy, T. B.; Chao, T.H.; Laane, J. J . Mol. Srrucr. 1972, 12, 427-449.
3352 The Journal of Physical Chemistry, Vol. 89, No. 15, 1985 brations unique to the keto form can account for its significantly greater relative entropy. The torsional frequencies are -40 cm-’ for reasonable model potential functions and are lower by a factor of -5 compared to those of the hydrogen-bonded enol form.” With the above assumptions the calculated entropy difference, AS(E-K) is -9-10 cal/(mol K) in good agreement with the experimental values obtained for all three molecules. This study has also determined the phase effects on thermodynamic parameters associated with the keto-enol equilibria. Interesting trends in Table I include the decrease in the relative enol-keto energy difference upon solvation as well as a slight but significant decrease in the relative entropy difference. Since polarity34-’Jand solvent packing are both important determinants of phase effects on equilibria, both factors are discussed qualitatively below. The keto and enol forms have appreciably different dipole moments, since the nature and relative orientation of the C - O bonds are significantly different in the two forms. The dipole moment of gaseous acetylacetone, measured from 323 to 573 K, is consistent with a total dipole moment of 3.00 D39for the enol form. For the keto form C=O bond dipoles of 2.2 D were assumed.@ From these values and the geometry reported in the electron diffraction study the total dipole moment of the keto tautomer is calculated to be -4.00 D. The gas-liquid change in the free energy difference between the two tautomers due to electrical effects may be calculated from reaction field theory by the equation:
where c is the dielectric constant for the medium, 9 is the solvent molecule radius, and p k and fie are the dipole moments of the keto and enol forms, respectively. The dielectric constant, e, for acetylacetone is 21.2 Using a solvent radius of 2.40 A, the calculated A(AG) is -3.3 kcal/mol. In solution solvent internal pressure will favor the form of the molecule which is ~ m a l l e r . ~ ~The . ’ ~ relative volumes of both tautomers were calculated by using a computer program written by H e r m a n r ~ , with ~ ~ geometrical parameters determined by electron diffraction and hard-sphere radii of 1.5 A for C and OU4* (34) See,for example: Abraham, R. J.; Bretschneider, E. In “Internal Rotation in Molecules”;ONilleThomas, W. J., Ed.; Wiley: New York, 1974; Chapter 13, pp 480-584. (35) Jorgensen, W. L. J . Phys. Chem. 1983,87, 5304-5314. (36) Chandler, D.; Pratt, L. R. J. Chem. Phys. 1976, 65, 2925-2940. (37) Pratt, L. R.;Chandler, D. J. Chem. Phys. 1978, 68, 4202-4212. (38) Hsu, C. S.;Pratt, L. R.;Chandler, D. J. Chem. Phys. 1978, 68, 4213-4217. (39) Zahn, C. T. Phys. Rev. 1933, 34, 570-574. (40) Cumper, C. W. Tetrahedron 1969, 25, 3131-3138. (41) Hermann, R. B. Program 225, Quantum Chemistry Program Exchange, Department of Chemistry, Indiana University, Bloomington, IN.
Folkendt et al. The calculated molecular volume of the enol form is 59.6 A3,and the volume of the keto form is 62.6 As. The corresponding enol-keto volume difference, AV(eno1-keto) is -3.0 cm3/mol. The enol form is more compact than the keto, primarily due to overla of the oxygen atoms. The nonbonded 0-0distance is 2.381 !f in the enol and 2.767 A in the keto tautomer. Forces which are purely steric in nature favor the enol form relative to the keto. The difference in AG for the ketc-enol equilibrium in the gas phase and in solutions due to steric effects may be calculated according to
where Pi is the internal pressure of the medium. With a AVof -3.0 cm3/mol and a liquid internal pressure of 3000 atm, and assuming that the gas is ideal (Pi = 0 atm), then the corresponding A(AG) is 207 cal/mol. Thus, the experimental result that the keto form is stabilized by -2 kcal/mol in condensed phases can be accounted for qualitatively on the basis of electrostatic effects which favor the form with the greater dipole moment. Steric effects favor the smaller enol form in the condensed phase, but this effect is insignificant in comparison to electrostatic factors. The Tl’s of the hydroxylic protons in the enol forms of the 8-dicarbonyls are also of interest. The symmetry of the H-bond potential function and the length of the T I are correlated. A significant contribution from tunnelling in the symmetric enols would account for this observation. The T I for acetylacetone’s hydroxylic proton is ca. 50 ms and is inversely dependent on temperature. Similar short T,’s for the hydroxylic protons have also been observed in the symmetric 8-decarbonyls 1,1,1,5,5,5hexafluoroacetylacetone, 3-methylacetylacetone, 3-ethylacetylacetone, 3-propylacetylacetone, and 1,s-diphenylacetylacetone. Conversely, the hydroxylic proton Ti's for methyl and ethyl acetoacetate, l,l,l-trifluoroacetylacetone,and l-phenylacetylacetone are ca. 500 ms. The values of T I for the other sites in all of the @-diketonesstudied are all 1-2 s. Further studies in this area are being conducted.
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Acknowledgment. We are pleased to acknowledge support from the National Science Foundation (Grant CHE82-10844 and CHE83-51698 (PYI)) and the National Institutes of Health (Grant PHSGM 29985). Registry No. AcCH2C(0)CH3,123-54-6; AcCH2C02Me,105-45-3; AcCH&O,Et, 14 1-97-9.
Supplementary Material Available: Tables of equilibrium constants (7 pages) for acetylacetone, ethyl acetoacetate, and methyl acetoacetate. Ordering information appears on the current masthead page. (42) Bondi, A. J . Phys. Chem. 1964, 68, 441-451.
