435
HINDERED ROTATION IN N-METHYLTHIOUREA
Hindered Rotation in N-Methylthioureal by A. S. Tompa, R. D. Barefoot, and E. Price Naval Ordnance Research, Naval Ordnance Station, Indian Head, Maryland
20640
(Received September 28, 1067)
A study of the barrier to rotation about the carbon-nitrogen bond in N-methylthiourea in pyridine, acetonitrile-&, trifluoroacetic acid, methanol, water, and deuterium oxide was performed by the nmr technique. In these solvents it was observed that below 5' there were two methyl doublets of unequal intensity and coupling constants. The coupling constant of the high-field doublet was 0.5 HZgreater than that of the low-field doublet and was assigned to the trans isomer. From the ratio of the trans to cis areas equilibrium constants were calculated and were found to be solvent and temperature dependent. From the coalescence of the methyl doublets approximate activation energies were calculated and found to increase slightly as the polarity of the solvent increased.
Nuclear magnetic resonance studies of the rotational H CH3 barrier about the central C-N bond in amides and / / HN thioamides have indicated that the energy barrier is CHSN greater than for a normal C-N bond because of electron \ \ delocalization.2-6 Furthermore, several authors have cscsreported that the rotational barriers in thioamides are +/ +// HN higher than for amides.4-* However, the determinaHN tion of kinetic parameters for the interconversion of \ \ geometrical isomers of amides and thioamides by nrnr H H spectroscopy has been a controversial s u b j e ~ t . ~ , 9 - ~ ~ trans cis Several authors have discussed and presented comparative results pertaining to the various nrnr methods These two isomers (cis and trans) show a chemical employed to obtain energy barriers to restricted rotashift difference between the methyl protons due to diftion about the C-K bond in amide^.^,^^^^ In spite of ferences in shielding by the sulfur atom. the errors involved in comparing results obtained by It was observed for N-methylthiourea in pyridine the various nrnr techniques, solvent and concentration 6' there were two methyl doublets of unthat below effectsseem apparent in these ~ysterns.'0-~~ In connection with a research program in our labo(1) Presented before the Division of Physical Chemistry a t the 153rd ratories, we have studied by the nmr technique the National Meeting of the American Chemical Society, Miami Beach, Fla., April 1967. behavior of N-methylthiourea (MTU)'O and N-methyl(2) C. C. Lin and J. D. Swalen, Rev. Mod. Phys., 31, 841 (1959). urea in a few nonpolar and polar solvents at various (3) H. S. Gutowsky and C. H. Holm, J. Chem. Phys., 25, 1228 temperatures in order to determine whether hindered (1956). rotation is present in these molecules and whether any (4) A. Lowenstein, A. Melera, P. Rigny, and W. Walter, J. Phus. Chem., 68, 1697 (1964). solvent interactions exist. The results and conclusions ( 5 ) R. C. Neuman, Jr., and L. B. Young, ibid., 69, 2670 (1965). of these studies are given below. (6) R. C. Neuman, Jr., D. N. Roark, and V. Jonas, J. Amer. Chem. There are two possible configurations of N-methylSoc., 89, 3412 (1967). (7) W.Walter and H. Maerten, Ann. Chem., 669, 66 (1963). thiourea due to restricted rotation about the C-N ( 8 ) W.Walter, G. Maeten, and H. Rose, ibid., 691, 25 (1966). bond
H +/
CH, +/
CHaN
HN
\\
7-
HN
\
H
trans
\\ cs/
HN
\
H
cis
(9) A. Allerhand and H. S. Gutowsky, J. Chem. Phys., 41, 2115 (1964). (10) E.Lustig and W. B. Monie, Anal. Chem., 38, 334R (1966). (11) A. Allerhand, H. S. Gutowsky, J. Jonas, and R. A. Meineer, J. Amer. Chem. Soc., 88, 3185 (1966). (12) M. T. Rogers and J. C. Woodbrey, J. Phys. Chem., 66, 640 (1962). (13) A. G. Whittaker and 8. Siegel, J. Chem. Phys., 42, 3320 (1965). (14) J. C. Woodbrey and M. T. Rogers, J. Amer. Chem. SOC.,84, 13 (1962). (16) E.Lustig, W. R. Benson, and N. Duy, J. Org. Chem., 32, 851 (1967). (16) Since the preparation of this article, T. H. Siddall and W. E. Stewart, J. Org. Chem., 32, 3261 (1967),have reported evidence of slow rotation of N-phenyl-N',N'-dimethylthiourea. Volume 73pNumber B February 1068
436
A. 8. TOMPA, R. D. BAREFOOT, AND E. PRICE
II
39%Trans
I
61% Cis
A
Figure 2. The nmr spectra of N-methyl protons in N-methylthiourea in HzO and DzO at 3".
+30.5O
A
Figure 1. The nmr spectra of N-methyl protons in N-methylthiourea in pyridine at different temperatures.
equal intensity and coupling constants (Figure 1). At 14" in pyridine, a broad singlet was observed which split into a doublet above 25". The doublet became very sharp as the temperature was raised to 61" and gave no indication that it would collapse to a singlet. However, in water the two doublets coalesced a t approximately 21" to give one doublet which collapsed to a singlet a t approximately 45". These results indicated to us that there are two kinetic processes occurring in the hydroxylic solvents: (1) hindered rotation about the C-;"\' bond and (2) proton exchange with the solvent. As shown in Figure 1, the high-field doublet had a coupling constant of 4.8 Hz, which was 0.5 Hz greater than that of the low-field doublet. The low-field doublet was assigned to the trans isomer and the high-field doublet to the cis isomer. This assignment is based on the following two factors. First, molecular models indicate that the trans isomer is expected to be less favored than the cis form because of steric interaction between the trans-methyl and the trans-hydrogen. Therefore, one would expect to find two unequal methyl doublets in the pmr spectrum. Second, the methyl protons cis to the thiocarbonyl group are expected to appear more upfield than the methyl protons on the opposite side because of anisotropic diamagnetic shielding by the thiocarbonyl group.l7JB Chemical shift and coupling The Journal of Physical Chemistry
constant data in various solvents and a t different temperatures are shown in Table I. I n deuterium oxide, the doublet due to spin-spin coupling disappeared because the protons on the nitrogen atoms were replaced by deuterium through proton exchange. Therefore, the methyl protons in Nmethylthiourea gave two singlets of unequal intensity for each isomer. The spectrum of the N-methyl protons of N-methylthiourea in deuterium oxide a t 3" is shown in Figure 2. These lines coalesced a t a temperature of approximately 29", whereas the two doublets in water coalesced a t approximately 21". These results give conclusive evidence that the unequal methyl doublets observed in water and the other solvents are due to the trans and cis isomere. From the ratio of the areas of the methyl doublets, equilibrium constants, K = [cis]/ [trans], were calculated in pyridine and deuterium oxide. Results are listed in Table 11. Between -30 and 5" in pyridine, the doublets were well resolved, but above 5" they began to overlap considerably and area ratios were difficult to obtain. I n deuterium oxide, the area ratio of cis to trans remained constant between -3 and 20" and then became equal as the singlets began to merge. On applying the method of least squares to the data in pyridine, we obtained the following results log K = -1.52
+ (426/T)
(1)
The standard error of fit equals h0.006 log unit, A H o = -1.95 kcal, AS" = -6.94 eu, and Ksas = 0.81. The only explanation that we offer concerning (17) It is well known that the carbonyl group exhibits anisotropic diamagnetic shielding. The anisotropy is such that protons located approximately in line with the double bond are shifted downfield, while those located approximately below or above the double bond are shifted upfield; cf. L. M. Jackman, "Applications of Nuclear Magnetic Resonance Spectroscopy in Organic Chemistry," Pergamon Press Inc., New York, N. Y., 1959, p 124. (18) H. W. Brown and D. P. Hollis, J. M o l . Spectrosc., 13, 305 (1964).
HINDERED ROTATION IN N-METHYLTHIOUREA
437
Table I : Nmr Chemical Shift Data for Methylthiourea in Various Solvents at 60 M R z
Pyridine
4.05
-21
-7 Methanol Acetonitrile-& Water Trifluoroacetic acid
1.57
- 30
0.94 0.80 1.11 0.33
2 -21 -24 3 - 11
4.1 4.4 4.2 4.0 4.2 4.4 4.3 1.0
4.6 4.8 4.6 4.6 4.8 4.9 5.0 2.0
2.88 2.97 2.98 2.92
3.10 3.18 3.23 3.15
...
...
2.68 2.82 2.91
2.87 2.95 2.96
8.25 8.15
9.08d 9.12
...
... ... ...
...
... ...
...
...
...
6.92
7.43d
Chemical shift, in ppm, relative to internal tetramethylsilane except in water where sodium 2,2-dimethyl-2-silapentane-5-sulfonate was used. The high-field methyl resonance has been assigned to the methyl protons that are cis to the sulfur atom (see the text). Average values of 4.3 0.1 and 4.8 i 0.2, respectively, for trans and cis isomers were obtained in the various solvents in the range The area ratios are 2.0 (NHn:NH) and of -30-5', with the exception of trifluoroacetic acid in which proton exchange occurs. 1.0 (CHs:(NH2 NH)).
+
Table I1 : Equilibrium Constants for trans-to-cis Interconversion of N-Methylthiourea in Pyridine and Deuterium Oxide OC
€a
8.5
Pyridine
1.55 1.46 1.37 1.28 1.19 1.10 1.01
-5 0 5 Deuterium Oxide
-3 2 17 20
Solvent
Pyridine
Temp,
-26 -20 - 15 - 10
Table 111: Solvent Effects on the Chemical Shiftb of the Methyl Protons in N-Methylthiourea
2.30 2.30 2.30 2.10
Area measurements were made with a planimeter.
the larger ratio of the cis to trans isomer in deuterium oxide is that some types of hydrogen-bonded complex stabilize the cis isomer preferentially. The only support that we have for such a complex is the observation that water gave approximately the same results as deuterium oxide, whereas methanol seemed to have shown higher values than pyridine and deuterated acetonitrile had lower values than pyridine. Gosavi, Agarwala, and Rao have reported a similar behavior of N-alkylthioureas in various solvents using infrared spectroscopy.19 From the separation between the two doublets, activation energies and frequency factors were calculated using the method of Gutowsky and The free energies of activation were computed according to the relationship given by Rogers and Woodbrey.12 Values for EA, A F j ) and log A ranged from 9 to 12 kcal/mol, 13.2 to 14.4 ltcal/mol, and 8 to 10, respectively, in going from pyridine to deuterium oxide.
Trifluoroacetic acid
Mol %
2.66 11.20 24.00 2.36 10 50 1.96 6.00 3.60 2.30 4.00 1.80 2.80 I
Acetone-&
20.7
Methanol Acetonitrile-&
32.6 37.5
Do0
78.0
A",b
17.2 14.5 12.5 4.5 4.2 5.5 5.0 11.3 11 .o 11.6 8.9 8.4
' Dielectric constants a t 25.0': "Handbook of Chemistry and Physics,'' 47th ed, Chemical Rubber Publishing Co., Cleveland, Ohio. * Chemical shift, in hertz, difference between methyl protons in the two isomers. Measurements were obtained in all solvents at -20 and -30' except in D20 at -3'. Inspection of the data in Table 111 reveals that the maximum value of the chemical shift, Avm, a t low temperatures strongly depends on the solvent and also on the concentration in pyridine. These values seem to be real and to be constant between - 10 and - 30°, with no indication of increasing as observed in some cases by Whittaker and Siegel.I3 These authors have indicated that changes in Av, (the chemical shift between the methyl groups nt which the internal rotation is negligible) with further lowering of the temperature arise because of the solvent or associative effects.
Experimental Section N-Methylthiourea (Aldrich) was recrystallized from absolute ethanol-ethyl acetate. Deuterium oxide (99.8%), acetonitrile-& (99.0%), and trifluoroacetic acid were obtained from KMR Specialties and were (19) R. K. Gosavi, U. Agarwala, and C. N. R. Rao, J. Amer. Chem. Soc., 89, 235 (1967). Volume YS, Number B Februarg 1969
438
LOUISWATTSCLARK
used without further purification. Pyridine (Mallinckrodt) and methanol (Matheson, spectroquality reagent) were freshly distilled before use. The nmr measurements were performed with a Varian DA-60-EL spectrometer equipped with a superstabilizer. The chemical shifts were measured with a precision of 0.1 Hz by placing side bands on both sides of the signal. The side-band frequency was measured with a Hewlett-Packard Model 522-B electronic frequency counter. The temperature was kept constant to within d~0.2"by the use of a Leeds and Korthrup
Azar H recorder-controller. The temperature was varied with dry nitrogen gas and the use of a Varian l\Iodel 4340 variable-temperature probe assembly and a hIodel V-433 1-THR spinning-sample dewar probe insert. Acknowledgment. This work was supported by the Foundat'ional Research Program of the Bureau of Naval Weapons. We also wish to thank Dr. James U. Lowe, Jr., for supplying us with a recrystallized sample of N-methylthiourea.
Further Studies on the Decarboxylation of n-Hexylmalonic Acid in Polar Solvents by Louis Watts Clark Department of Chemistry, Western Caroldna University, CulZowhee, North Carolina R87X3
(Received July 8 1 , 1 9 6 8 )
Rate constants and activation parameters are reported for the decarboxylation of n-hexylmalonic acid in 12 polar liquids-1 ,3-propanediol, lf2-propanediol, glycerol, 2,3-butanediol, 1-decanol, 1-octanol, 1-hexanol, K,S-dimethylaniline, o-toluidine, N-methylaniline, o-ethylaniline, and aniline. I t is observed that the inductive effect of the n-hexyl group of the substrate causes an inversion of the order of increasing activation energy with increasing molecular weight of the members of the various homologous groups (glycols, alcohols, and amines). The average change in the enthalpy of activation per methylene group for the decarboxylation of n-hexylmalonic acid in the amines is compared with corresponding values previously reported for five other acids-malonic acid, oxanilic acid, oxalic acid, ciiinamalnialonic acid, and benzylmalonic acid. The values obtained for all the compounds turn out to be exact multiples of the absolute magnitude of the value for malonic acid. A plot of A H $ us. As$ for the decarboxylation of n-hexylmalonic acid in each homologous group is linear, each line having a slope of 378'K or 105°C-the same temperature as the melting point of n-hexylmalonic acid. An interesting comparison is made between the data for the decarboxylation of n-hexylmalonic acid in alcohols and that for the solvolysis of methyl p-toluenesulfonate in alcohols reported by Hyne and Robertson. The melting point of p-toluenesulfonic acid is 105O-the same as the melting point of n-hexylmalonic acid. The isokinetic temperature for the solvolysis of the ester in alcohols is also 105O-the same as that for the decarboxylation of n-hexylmalonic acid in alcohols. Finally, the average change of the enthalpy of activation per methylene group on going from one alcohol to another is exactly the same for both reactions, namely, 947 cal.
Justification for continued investigation of the kinetics of the decarboxylation reaction lies in the potential of such studies to throw light on the mechanism and energetics of liquid-phase reactions. The rate-determining step in the decarboxylation of malonic acid in the presence of quinoline appears to be a nucleophilic attack by quinoline on the hydrogen-bond donating carboxyl group of malonic acid.' The electrophile is cleaved and the nucleophile escapes unscathed from
mechanism, the study of this reaction cannot fail to yield greater understanding of the properties of chemical species in general. n-Hexylmalonic acid is a suitable compound for use in studying this reaction because it is easily purified and readily undergoes reaction in all sorts of polar liquids at moderate temperatures. The equation of the reaction is CHaCH&H&H2CI-I2CH2CH (COOH) 2 COz
4
+ CHaCH2CH2CH2CHzCHzCH2COOH
The decarboxylation of n-hexylmalonic acid has been (1) G.Fraenkel, R . L. Belford, and I?. E. Yankwich, J . Amer. Chem. Soc., 76, 15 (1954). (2)L. W.Clark, J . Phys. Chem., 71,302 (1967). and previous papers
the collision. This mechanism apparently obtains in a large variety of decarboxylation reactions.2 There appears to be no limit to the number of acids able l o undergo cleavage nor to the number of nucleophiles capable of effecting it. Because of this generality of The Journal of Phgsical Chemistry
in this series. [A reviewer has suggested that the alternative mechanism of nucleophilic attack by quinoline on the hydrogen-bond accepting carboxyl group i s not ruled out by the data of ref 1 and 2. In addition, this reviewer stated, catalysis by this alternative mechanism is easily understood and is consistent with the effects of substituents on the decarboxylation of @-keto and unsaturated acids-cf. D. B. Bigley and J. 0. Thurman, J . Chem. SOC.,E , 436 (1968)I.
