A STUDY OF TAUTOMERISM IN CYCLIC Q-DIKETONES BY PROTON MAGNETIC RESONANCE

NATSUI~O CYR AND LEONARD W. REEVES CJaenzistry Departnzent, University of British Colz~nzbia, Vancouaer, British Colz~nabia

Received May 25, 1965

ABSTRACT

The Iteto-en01 equilibrium of cyclohexane-l,3-dione in chloroform is best interpreted from proton resonance ~neasurements as

K I R z en01 dilner 2-en01 monomer 2-lrcto monomer.

K1 and Kz may be separately determined from chemical shift measurements of the enol-OH proton and intensity measurements of peaks assigned to keto and en01 forms. K1 and R? are satisfactorily independent of concentrations except in very dilute solutions where intensity measurements become unreliable. The overall equilibrium constant K = K1 X K ? can be obtained for the same ~nolecule in acetonitrile solutions where the en01 monomer form is in very low concentration. 5,5'-Di~nethylcyclohexane-1,3-dione in chloroform has less en01 form than the unsubstituted molecule. The enthalpy change associated with 'K' for cyclohexane-1,3- dione in chloroform is 2.05 f 0.5 kcal mole-'.

INTRODUCTION

I t is well known that /3-diltetones forin en01 isomers and that the extent of enolization is controlled mainly by the conjugation of C=O and C=C double bonds and the strong intramolecular hydrogen bond formed (1, 2). The equilibrium keto en01 is affected by concentration and solvents as well as temperature (3). Proton magnetic resonance studies of such tautomeric systems generally reveal separate pealts for each tautoiner thus enabling a simple intensity measurement to give the equilibrium constant (3, 4). There are several interesting features of proton magnetic resonance studies of tautomeric systems (5-13) including spin-spin coupling to the enol-01-1 proton (11) and multi-site proton exchange processes which can be determined by double resonance methods (14, 15).

The present study is directed towards the six-membered cyclic /3-diketones where the formation of an intramolecular hydrogen bond is excluded. Cyclohexane-1,3-dione and 5,s'-dimethylcyclohexane-1,3-dione exist both in the solid state and concentrated solutions mostly as enolic n~olecules (16-18). A large diiner unit has been suggested (IG) to provide the strong hydrogen bond which stabilizes the en01 form. The study of the tautomeric equilibrium in these n~olecules is confined to chloroform and acetonitrile solutions where the solubility is appreciable.

EXPERIMENTAL

Commercially available samples (L. Light and Co. and Eastman Icodak Co.) of cyclohexane-l,3-dione and 5,5'-dimethylcyclohexane-1,3-dione were purified by recrystallization from benzene. Spectroscopic grade chloroform was shaken with concentrated sulfuric acid, washed with distilled water, and fractionally distilled to remove added ethanol. I t was stored over calcium chloride in a dark place. Reagent grade acetonitrile was fractionally distilled and kept over calcium chloride. All concentratioils were obtained by weighing solute and solution. The solutions were sealed in vacuum in nuclear magnetic resonance (n.m.r.) sxrnple tubes after degassing by successive freeze, pump, thaw procedures. A small amount of magnesium perchlorate for the chloroform solutions and sodium sulfate for the acetonitrile solutions was added to each i1.m.r. sample tube to keep solutions dry throughout measurements. A Varian A60 spectro~neter was used for measurements with accurate calibrations provided by a Hewlett-Paclrard Wide Range Audio oscillator standardized a t each measure~neilt by a model 5223 electronic frequency counter. Higher temperatures in the A60 spectro~neter were obtained by using a V6057 variable temperature attachment and temperature measurement achieved with the calibrated separation of methylene and -OH proton resonances in ethylene

Canadian Journal of Chemistry. Volume 43 (1965)

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3058 CANADIAN JOURNAL O F CHEMISTRY. VOL. 43, 1965

glycol. A small amount of tetramethylsilane (TMS) served as an internal reference for chemical shifts. The intensities of respective peaks in en01 and keto forms were measured a t the slowest sweep rate: 500 s on the narrowest sweep width, 50 c.p.s., and with the spectrum amplitude not exceeding the linear response region of the instrument. The absorption signals were carefully traced onto good quality bond paper and the cutouts weighed.

The signals used for intensity measurements are: Cyclohexane-1,3-dione in CHCl,

enol-OH (7 in structure I) Iceto-CM2 between the carbonyls (2 in structure 11).

Cyclohexane-1,3-dione in CHICN enol-olefinic proton (2 in structure I ) keto-CH2 between the carbonyls ('2 in structure 11).

5,5'-Dimethylcyclohexane-l,3-dione 4 and 6 signals for both the keto and the enol.

Chem~cal sh~f t . c.p.s. from TMS at 60 M C

Chemical shift , cp.s, from TMS at 60 Mc

FIG. 1. . (a ) Spectrum of cyclohexane-1,3-dione in chloroform 16.0 mole %. (b ) Spectrum of cyclohexane- 1,3-dione In chloroform 1.05 mole 70.

RESULTS AND DISCUSSION

A . Eq~iilibrizirn Constants The spectrum of cyclohexane-1,3-dione a t the solubility limit of 1G.O inole % in chloro-

form is shon-n as Fig. la . The pealcs can all be reliably assigned to the en01 form. As ill structure I we have enolic-OF1 a t -700c.p.s. (-ll.G7 p.p.m.), olefinic C-H

a t position C2 -330 c.p.s. (-5.5 p.p.m.), aliphatic -CH2- protons giving a colnplex spectrulll of which peaks centered a t -120 C.P.S. (-2 p.p.m.) can be attributed to 5C

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CYR AND REEVES: TAUTOMERISM IN CYCLIC 8-DIKETONES 3059

protons. In Fig. lb a t 1.05 mole % in CHC13 signals due to the Iceto form (structure 11) appear. These have been shaded to make them more distinct in the figure.

Peaks a t -207 c.p.s. (-3.45 p.p.m.) are assigned to the C2 methylene protons and the C4, C5, and C6 protons of I1 are in a region - 155 c.p.s. (-2.59 p.p.m.). A noticeable feature of Fig. lb is the change in chemical shift of the enol-013 proton to -580 c.p.s. (-9.66 p.p.m.) from Fig. l a . The large change in chemical shift of the enol-OH proton with concentration is indicative of the rupture of hydrogen bonds (19) as the en01 form is diluted. The chemical shift of the enol-OH proton has been measured accurately a t all available concentrations and the variation can be described by a linear dependence on l / d C . Figure 2 shows a plot of chemical shift from T M S versus the inverse square root of concentration in (moles 1)-lJ2. The chemical shift a t infinite concentration is -740 c.p.s. (-12.32 p.p.m.) and must correspond to the completely hydrogen-bonded en01 species, most reasonably represented as a dimer. (Structure III.)*

FIG. 2. Chemical shift of enol-OH proton as a function of C-"2 (C = concentration in moles/l). FIG. 3. Dependence of log K for cyclohexane-1,3-dione in chloroform on inverse absolute temperature.

This dimer \vould account for the very low field shift of the -OH proton and is typical of the shifts observed for intramolecular hydrogen-bonded enols (4, 5). The large con- centration dependence of the enolic-OH chemical shift must result from dissociation of

*The 0-H . . 0 lines are approxinzately i n the plane of the cyclohexane ring because the lone pairs on the oxygen determines the strongest hydrogen bonds i n this direction. The olejinic hydrogens at C2 are not close enough to interact with each other.

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3060 CANADIAN JOURNAL OF CI-IEMISTRY. VOL. 43, 1965

these diiners and the appearance of mono~neric en01 inolecules in solution. The equilibria in chloroforin solutions can therefore be described by

K1 Kz en01 diiner S en01 monomer $ keto monomer Co(1 - or) 2or CD 2(C - CD).

Concentrations a t equilibiruin in terins of initial concentrations of undissociated forms are given below the equation above. C = total concentration of cyclohexane-1,3-dione expressed in units of diiner (moles I-'); CD = concentration of en01 form considered coinpletely as dimeric forin (moles I-') ; a = degree of dissociation of the en01 form into monomers.

TWO equilibriu~n constants can be defined as above in terms of these concentrations.

The value of CD is determined from a knowledge of the signal areas for en01 and keto forms

[31 C, = C/(I + ( 1 / 2 ~ ) )

where R = [enol]/[keto]. The degree of dissociation 'a' \vill be obtained from the observed hydroxyl proton chemical shift

XD and Xxx are the respective fractions of en01 molecules in dimer and monomer form. 6, and are chemical shifts associated with the enol-OH proton in dimer and monomer forms of the en01 molecule. 6obs is the observed chemical shift a t some arbitrary con- centration. Froin Fig. 2 the value of 6, may be taken as -740 c.p.s. Thus

The only unknokvn is ' 6 ~ ' which may be chosen so as to give constant values of K 1 and K z a t all concentrations. The best value is -300 c.p.s. (-5 p.p.m.) giving a change in chemical shift due to hydrogen bond formation of 7.32 p.p.m. This is in close agreement with previous values suggested for a strong -0-H --- 0 hydrogen bond (20, 21). The equilibrium constants obtained are listed in colulnns 8 and 9 of Table I . The values of K1 are satisfactorily constant for all concentrations but values of K 2 vary between 1.13 and 0.42 over the concentration range studied. There is a reasonably constant value of K 2 = 0.55 =t 0.01 a t the higher concentrations of cyclohexane-1,3-dione. The possible factors likely to cause errors are: (a) Impurities contained in the chloroform which contain -OH groups and enhance the enol-OH peak intensity a t low concentration. (b) Low signal to noise and possible saturation of the resonances a t lo\v concentrations. This \vould cause broadening of signals and loss of intensity in the noise. If relaxation times for the resonances chosen for intensity measurements are different there could be a differential factor (22). (c) The solubility of n~agnesiun~ perchlorate in chloroform is small a t room temperature. However, in ketonic solvents such as acetone it is quite large. In the presence of a large quantity of cyclohexane-1,3-dione the increase of the solubility of the salt could cause soine sinall error as well.

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CYR A N D REEVES: TAUTOMERISM I N CYCLIC 8-DIKETONES 3061

TABLE I Equi l ibr i~~m constant data for cyclohexane-1,3-dione in chloroform (C = concentration in moles 1-1, -6,b, = chemical shift of -OH proton from TMS in c.p.s., Chr = concentration of monomer en01 form moles

I-'. All other symbols are defined in equations in the text)

c CD CH ~ ( C - C D ) K l ( x 10) R ( X 10) -sobs a ( X lo2) ( X lo2) ( X lo2) Kz

Average 7 . 3 M o F 2

The purified chloroform was run as a blank sample and no impurities were found at the highest sensitivity of the spectrometer. The effect of (b) above would be most serious in measurements of ICz a t low concentration where the ratio 'R' occurs in both numerator and denominator of eq. [2]. I t is reasonable thus to neglect values of Kz below 1.4 moles/l.

The equilibrium constants obtained for acetonitrile solutions of cyclohexane-1'3-dione are obtained from direct n~easurements of intensities of keto and en01 peaks assuming negligible monomer en01 form is present. The equilibrium constant 'K' may be defined

The values obtained for I< over a threefold increase in concentration are satisfactorily constant as shown in Table 11. Corresponding values of K = K1 X KZ2 for 5'5'-dimethyl- cyclohexane in chloroform solutions were necessarily limited to the concentration range O.G7 to 0.87 moles/l by the solubility and spectrometer sensitivity. A constant value of I< = 5 .5 f 0.4 X lop2 was, however, found in this limited range. This compares with I< = 2.21 f 0.2 X for the cyclohexane-1,3-dione itself in the same solvent from Table I. The lteto form is thus favored slightly by the 5'5'-dimethyl substitution.

TABLE 11

Overall equilibrium constant K for keto/enol system cyclohexane- 1,3-dione in acetonitrile solutions (symbols as defined in the text)

C (X 10) R CH ( X 10) Cn ( X 10) K ( X 10)

0.62 0.568 0 .29 0 .33 1 .O 0.95 0.699 0 .38 0.57 1 . 0

B. Temperature Dependence of 'K' The overall equilibrium constant 'K' for cyclohexane-1'3-dione in chloroform a t a

concentration of 0.47 mole/l was measured as a function of temperature. A t this con- centration, dissociation into lnonorner en01 molecules is small and 'K' can be measured fro111 peak intensities alone. The enthalpy change associated with 'K' is derived as 2.05 =t

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3062 CANADIAN JOURNAL O F CHEMISTRY. VOL. 43, 1966

0.5 kcal mole-I, from the slope of Fig. 3. This can be compared with the value 2.7 kcal mole-' obtained by Reeves (3) for pure acetylacetone.

ACKNOWLEDGMENTS

One of us (L. W. R.) is grateful to the National Research Council of Canada and the Petroleum Research Fund of the American Chemical Society for the financial support of this ~vork.

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10. G. 0. DUDEK and R. M. HOLM. J. Am. Chem. Soc. 84, 2691 (1962). 11. S. FORSEN and M. NILSSON. A r k ~ v Kemi, 19, 569 (1962). 12. G. 0. DUDEK. J. Am. Chem. Soc. 85, 694 (1963). 14. S. FORSEN and R. A. HOFFMAN. J . Chem. Phys. 39, 2892 (1963). 15. S. FORSEN and R. A. HOFFMANN. J. Chem. Phys. 40, 1189 (1964). 16. K. KODERA. Yakugaku Zasshi, 80, 1267 (1960). 17. C. DEVAL and J. LECOMPTE. Compt. Rend. 254, 36 (1962). 18. C. L. ANGELL and R. L. WERNER. Australian J. Chem. 6, 294 (1953). 19. L. W. REEVES and W. G. SCHNEIDER. Trans. Faraday Soc. 54, 314 (1958). 20. L. W. REEVES, E. A. ALLAN, and K. 0. STR@XME. Can. J. Chem. 38,1249 (1960). 21. E. A. ALLAN and L. W. REEVES. J. Phys. Chem: 67, 591 (1963). 22. N.M.R. and E.P.R. Spectroscopy. Varlan Associates. Pergamon Press, London. 1960. Chap. 8.

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