JAN SANDSTR~M
2318
Calculation of the latter is based on the density of solid HMT and although the comparison is in this case a qualitative one, it appears plausible to assume that an “excess” increase in volume takes place when 1 mole of HMT is dissolved in a large amount of liquid water. The negative slopes of the approximate VZ m relations (Figure 2) lead us to think that an increased ice-likeness of water in the HMT solutions is promoted by the solute. This view appears to be in agreement also with some features exhibited by the near-infrared spectrum of water in HMT solution (see Figure 4) which suggest
that HMT affects water structure in a way formally equivalent to a decrease in temperature.” The attempted description of the HMT-H20 system is of course still highly speculative, but one according to which the results reported here find a consistent though qualitative interpretation.
Aclcnowkdgments. The authors wish to thank Professor Alfonso M. Liquori for his helpful advices during the course of this work and Dr. Donald G. Miller for critically reading the manuscript. (11) K.Buije and G. R. Choppin, J. Chem.Phys., 39, 2035 (1963)
Barriers to Internal Rotation in Thioamides. Experimental Results and Molecular Orbital Calculations
by Jan Sandstrom Department of Chemietry, u n ~rrsita, . of h n Lun., Sweden Accepted and Transmitted by The Faraday Society (November 50,1966)
The influence of the substituent R in RCSN(CH& on the barrier to the rotation of the dimethylamino group has been studied by nuclear magnetic resonance. The Arrhenius parameters have been determined together with the AF*, A H f , and AS* values and the effects of solvent and concentration on these have been studied. The AF* values show a very qualitative correlation with C-N ?r-bond orders calculated by a modified w method and a better correlation with the loss in 1-electron energy, AE,, which occurs when the dimethylamino group is rotated out of conjugation. This quantity also gives correct relations between the barriers of thioamides, amides, and amidinium ions, and it allows a crude prediction of the barriers in amidines and enamines.
Introduction It has been observed by Luttringhaus,
et al.,1 that
the carbonyl group interacts more strongly than the thiocarbonyl group with weakly electron-donating groups, whereas the reverse is tJ” with strongly ekcThis is in agreement with the tron-donating- _groups. opinion expressed by several authors2 that thioamides are more than the corresponding amides. As has been shown by Loewenstein, et The Journal of Physicol Chemtktw
one consequence of this is a considerably higher barrier to internal rotation in N,N-dialkylthioformamides (1) (a) A. Lattringhaus and J. Grohmann, 2. Naturforsch., lob, 365
(1955); (b) A. Luttringhaus, R. Mecke, R. Mecke, and J. Grohmann in “Elektronentheorie den Homiipolaren Bindung,” AkademieVerlag, Berlin, 1956, p 152. (2) (a) C. M. Lee and W. D. R u d e r , J. Ow.Chem.,27,2052 (1962); (b) K.A. Jensen, Acta Chem. Scand., 17, 551 (1963). (3) A. Loewenstein, A. Melera, P. Rigny, and W. Walter, J . Phys. chem., a,1697 (ISM).
BARRIERS TO INTERNAL ROTATION IN THIOAMIDES
2319
than in dimethylformamide, as measured by standard nrnr technique. This observation is supported by the data for N,K-dimethylthioacetamide and N,Ndimethylacetamide obtained in the same way by Neuman and Young.' However, the Arrhenius activation energies reported by these authors are considerably higher than earlier values, probably because of the very polar solvent used (formamide). Quite recently, Walter, et ~ l .have , ~ obtained in a very elegant way the barrier in N-methyl-N-benzylthioformamide by studying the temperature dependence of the rate of cis-trans interconversion, as obtained from precise integration of the spectra of the two rotamers. Their data are in good agreement with those reported by Loewenstein, et aL3 It is apparent from the barriers for N,N-dimethylamides reported by Rogers and Woodbreys that the substituent on the carbon atom has a strong influence on the height of the barrier. This was interpreted as an effect on the bond order of the carbon-nitrogen bond. It has been found that the ultraviolet spectra' and also the polarities* of a wide range of thiocarbonyl compounds can be satisfactorily described by an LCAOMO method of modified w type and it was considered as a matter of interest to test this method on the barriers to internal rotration of a series of substituted N,Ndimethylthioamides (I), where the groups R have electron-a t trac t ing or electron-donating properties.
quirements and it has an AA'BB' spectrum with sharp lines, which is convenient for checking the resolution of the spectrometer. Recording of the N m r Spectra. The spectra were obtained a t 60 Mc/sec with a Varian A-60 spectrometer equipped with a Varian V-6031 variable-temperature probe and temperature controller. The spectra were recorded with a sweep width of 1 or 2 cps/cm, depending on the peak separation, and a t a sweep rate of 0.1 or 0.2 cps/sec. The amplitude of the radiofrequency field was kept well below the level where saturation effects could be observed. At each temperature five to ten spectra were recorded with both upfield and downfield direction of the recorder. Where a reference was reqwed, the chemical shifts were measured against hexaniethyldisiloxane as an internal standard. Its resonance falls 0.055 ppm downfield from tetramethylsilane. Temperature Measurement. For determining the temperature of the sample a narrow concentric capillary, containing ethylene glycol or acidified methanol, was inserted into the sample tube and the shift between the OH and CH signals was recorded simultaneously with the spectrum of the sample. The shifts were calibrated using the ordinary Varian ethylene glycol and methanol samples. In this way an accuracy of about h0.4" could be obtained. In each run the spectra were recorded at 10-12 different temperatures. This method failed for N,N-dimethyl-N'-acetylthiourea (I, R = CHaCONH) and ethyl N,N-dimethylthiooxamate (I, R = EtOCO) because of overlap of sample and temperature standard signals. In these cases the conventional method had to be used.
S-
S
'R--C
// \ N-CH3 /
R-C
/ \+
/
N-CHI
CH3 CH3 Ia Ib Experimental Methods Materials. The thioamides were prepared by literamethods' The 'lid were purified by distillation, followed by recrystallization and, in some cases, by vacuum sublimation. The only noncrystalline compound (I, R = CH,O) was purified by fractional vacuum distillation. The question of finding a suitable solvent is rather difficult, since it should remain liquid and have good dissolving properties in the temperature range -" to +2000 and there should be little interaction between solvent and solute molecules. Finally, o-dichlorobenzene was chosen (molar fraction of solute 0.333) following the example Of LoewensteinJ et is not as as should have been desired, but it fulfills the other re-
Calculation of Thermodynamic Parameters The calculation of the rate constants for the rotation around the carbon-nitrogen bond was performed according to the ratio method described in detail by Rogers and Woodhrey.6 Quite recently, Allerhand, et aL,g have critically examined the different methods for obtaining chemical-exchange rates from nmr data. From their results it is possible to estimate the magnitude of the systematic errors which are introduced by use of the ratio approximat'ion. It appears (4) R. C. Neuman, Jr., and L. B. Young, J . Phys. Chem., 69, 2570 (1965). (5) w . Walter, G* Maefien, and H. Rose, Ann. Chem., 691, 25 (1966). (6) M. T. Rogers and J. C . Woodbrey, J . Phya. Chem., 66, 540 (1962). (7) (a) M. J. Janssen, Rec. Tmv. Cham., 79, 1066 (1960); (b) J. Sandstrom, A& Chem. Scund., 20, 689 (1966), and earlier papers. (8) M. J. Janssen and J. Sandstrom, Tetrahedron, 20, 2339 (1964). (9) A. Allerhand, H. 9. Gutowsky, J. Jonas, and R. A. Meinzer, J . Am. c h m . SOC., 88, 3186 (1966).
Volume 71. N u m k 7 June 1987
JANSANDSTR~M
2320
that in most cases the error in 7 is less than 5%, but in one case, uiz., 1, R = C, the peak separation in the absence of exchange is only 4.0 cps and the parameters obtained for this compound must be regarded as somewhat less reliable. In all cases the error works in the same direction, ie., it will tend to decrease the slope of the log k us. 1 / T plot and therefore it is not expected to affect the correlation in a very serious way. Wittaker and SiegeIlO have observed that the methyl doublet separation in N,N-dimethylformamide in a variety of solvents is considerably sensitive to the temperature even at temperatures where the internal rotation must be quite slow. In some cases a variation of about 12% over a temperature range of 40" was observed. This was ascribed to a temperature-dependent dipoledipole association which affects the anisotropy of the carbonyl group. With one exception, this effect is rather small in the systems investigated here. From the temperature where exchange effects could first be observed and 10-20" or more downward, the doublet separation varied by less than 2%. One reason for this may be that solventrsolute interactions are more important and less temperature sensitive than solutesolute interactions. The only exception is ethyl N,N-dimethylthiooxamate (I, R = EtOCO), for which an increase in peak separation is observed with increasing temperature up to a ratio of about 5. This may be due to a temperature-dependent change in the ratio between cis and trans conformers. For this compound the largest observed peak separation was used for avo, but it is obvious that this complication renders the thermodynamic parameters for this compound somewhat uncertain. The rate constants and temperatures were fitted by the method of least squares to a straight line according to eq 1 and the other activation parameters were obtained by use of eq 24." The transmission coefficient logk =
E, - 2.303RT ~
+ log A
AF* = -2.303RT log
AH*
=
E,
As* = AH*
hk ~
kTx
- RT - AF* T
(1)
(3)
(4)
n is as usual assumed to be unity. The results are shown in Table I. The errors in E, and log A are standard deviations from the least-squares plot and the errors in AH* and AF* are assumed to be the same. Thus, the errors in a*are obtained by doubling the deviations in E, and dividing by the temperature. The
The Journal of Physicrrl Chemktru
activation parameters show a slight temperature dependence and the values in Table I refer to the coalescence temperature. In the spectrum of N,N-dimethylthioacetamide (I, R = CHa) a coupling with J = 0.6 cps is observed b e tween the C methyl group and one of the N methyl groups.12 This introduces an extra systematic error when the simple ratio method is used for calculation of rate constants, but an examination of the analysis in ref 9 shows that this error is only of the order of 1-2'3, in the present case. For tetramethylthiourea only one sharp signal waa observed down to -60". This can be due to an accidental overlap of the two methyl signals, as is observed for dimethylthioformamide in o-dichlorobeneene in the molar ratio 1 :2, but it seems less likely since the same result is obtained in different solvents and at different concentrations. A low barrier is an explanation which is more in line with the general result of this investigation. Molecular Orbital Calculations The calculations have been performed by a modified w method, in which Coulomb integrals are corrected according to n-electron charges and resonance integrals according to bond orders in an iterative procedure. It was found8 that in order to reproduce with this method the interaction of electron-donating substituents with carbonyl and thiocarbonyl groups, as manifested in dipole moments, a low-resonance integral for the thiocarbonyl group and an element of electronegativity for the sulfur atom were necessary. As parameters for substituent atoms and bonds, standard values were tried, but they were mostly found to give too large dipole moments. When the resonance integrals were scaled down to 0.67-0.75 of the original values, reasonable figures for the dipole moment differences between corresponding carbonyl and thiocarbonyl compounds were found. The parameters thus obtained (set 4 in ref 8) are shown in Table I1 and the calculated n-bond orders and ?r-electron energies in Table 111. Other parameter sets with w = 1.0 or 1.4 and with larger l c c ~or smaller hx values have also been tried and found to give qualitatively similar results in the barrier calculations, as long as the relation between the parameters for the carbonyl and thiocarbonyl group is the same as here. (10) A. G. Wittaker and S. Siegel, J. C h . Phys., 42, 3320 (1965); 43, 1575 (1965).
(11) A. A. Frost and R. G. Pearson in "Kinetics and Mechanism," 2nd ed, John Wiley and Sona, Inc., New York, N. Y., 1961, p 98. (12) R. C. Neuman, Jr., and L. B. Young, J. Phys. Chem., 69, 1777 (1965).
2321
BARRIERS TO INTERNAL ROTATION IN THIOAMIDES
Table I
R
Molar fraction of thioamide
Solvent
ODCu Neat ODC ODC ODC ODC ODC ODC CCh CHCh (CH3)GHOR ODC ODC ODC ODC ODC 0
0.333 1
0.82 0.64 0.50 0.333 0.20 0.333 0.333 0.333 0.20 0.333 0.333 0.333 0.333 0.333
*,
VB,
opa
Ea, kcal/mole
log A
kcaf/mole
kcal/rnole
182.0 183.6 180.8 177.6 175.3 172.0 168.8 189.4 200.0 199.2 188.9 191.6 167.6 189.6 186.7 182.5
21.0 f 0.3 16.4 f 0.2 16.0 f 0.5 16.7 f 0.3 17.0 f 0.3 17.3 f 0.2 17.6 f 0.4 14.0 f 0 . 4 11.7 f 0 . 4 11.5 f 0 . 5 11.2 f 0.6 13.7 zt 0.4 19.1 f 0 . 5 23.2 f 0 . 5 20.8 f 0.8 17.1 f 1.0
12.6f0.2 12.0 f 0.2 11.7 f 0.3 12'.2 f 0.2 12.4 f 0.2 12.6 f 0 . 2 12.8 f 0 . 3 11.6 f 0 . 3 9.9 zt0.3 9 . 6 f 0.4 9 . 6 f 0.4 9.4zt0.2 13.3 f 0.4 12.8f0.2 11.7 f 0 . 5 13.5 f 0 . 8
21.6f0.3 17.9 f 0.2 17.8 f 0.5 17.8 f 0.3 17.7 f 0 . 3 17.7 f 0.2 17.7 f 0 . 4 15.6 A 0.4 15.6 A 0.4 15.9 zt 0.5 15.7 f 0.6 19.1f0.4 18.4 f 0.5 23.4f0.5 23.4 f 0 . 8 16.111.0
20.2f0.3 15.8 f 0 . 2 15.3 f 0 . 5 16.0 f 0 . 3 16.3 f 0.3 16.6 f 0.2 16.9 f 0.4 13.4 f 0 . 4 11.12~ 0.4 10.9 f 0.5 10.6 f 0 . 6 13.0f0.4 18.4 f 0 . 5 22.3 f 0 . 5 19.9 f 0 . 8 1 6 . 4 1 1.0
196.5 195.5 194.0 192.7 191.6 190.1 187.8 203.2 207.0 208.8 201.4 195.6 199.5 201.6 193.1 195.6
AF
AH
AS
*,
BU
-3.5f1.5 -6.2 f 1.2 -7.3 f 2.7 -5.2 f 1.8 - 4 . 1 f 1.8 -3.2 f 1 . 2 -2.5 f 2 . 3 -7.4 f 2.7 -15.2 f 2.7 -16.4f3.3 -16.7 f 4 . 0 -17.5f2.3 -0.1 f 2.7 -2.6 f 2 . 3 -7.9 f 3.7 $1.2 f 6.5
To OK
9
414.1 340.9 341.0 342.6 342.7 343.4 345.0 302.2 294.3 302.4 303.0 350.2 365.0 443.9 432.6 309.7
ODC, o-dichlorobenzene.
Table I11 : Bond Orders and ?r-Electron Energies (in Units of 8 )
Table 11" Bond
Atom (X)
C N N cj
is
s s
C1 0
*,
PA,
CPB
w =
1.4.
'
hxb
(C-W
kcxC
0 0.5 1.5 1.0 2.5 0.5 1.0 3.0
c-C C=N C-N C 4 C-S C-Cl
0.75 0.9 0.8 0.8 0.4 0.4
CYX =
cuc
R
H H RO RO RS
Rs
+ ~ X ~ C C . BCX = kcx8cc.
Discussion Eflect of Solvent and Concentration. The influence of the solvent on the thermodynamic parameters was studied only with methyl N,N-dimethyldithiocarbamate (I, R = CHBS). It was expected that hydrogen bonding should increase the barrier by stabilizing the polar structure Ib more than Ia. However, no such effect was observed within the limits of experimental accuracy. The AF* and A S * values are almost the same in carbon tetrachloride, chloroform, and 2-propanol. On the other hand, in o-dichlorobenzene, AF* has almost the same value, whereas A S S is about 8 eu higher. This corresponds to a greater loss of order on rotation in the latter case and the effect can be rationalized in terms of complex formation between the aromatic solvent molecules and the thioamide molecules. Such complexes have been proposed by
c1 c1 CHFCH CH-CH Ph Ph N=C N=C EtOCO EtOCO AcNH AcNH NR2
NRa
X
PCN
0 S 0 S 0 S 0 S 0 S 0 S
0.422 0.455 0.402 0.436 0.408 0.438 0.419 0.452 0.393 0.412 0.395 0,415 0.408 0.434 0.404 0.428 0,392 0.421 0.387 0.417
0 S 0 S 0 S 0 S
AEn
6.962 5.342 12.248 10.656 9.100 7.511 13.017 11.402 9.408 7.868 15.156 13.614 11.717 10.176 15.933 14.390 14.255 12.695 10.390 8.830
6.430 4.706 11.778 10.106 8.616 6.956 12.498 10.784 8.911 7.302 14.659 13.046 11.193 9.551 15.419 13.786 13.804 12.181 9.962 8.342
0.532 0.636 0.470 0.550 0.484 0.555 0.519 0.618 0.497 0.566 0.497 0.568 0.524 0.625 0.514 0.604 0.451 0.514 0.428 0.488
Hatton and Richards13to account for the shifts of the N-methyl signals in the nmr spectra of N,N-dimethylamides on dilution with benzene and other aromatic solvents. Similar solvent shifts have previously been and, in reported for N,N-dialkylthioformamidesS~6 connection with the present work, have also been ob(13) J. V. Hatton and R. E. Richards, Mol. Phys., 3, 253 (1960).
Volums 71, Number 7 June 1067
JANSANDSTROM
2322
served for a number of other thioamides. The results indicate, in agreement with the observations of Neuman and Young,12 that the magnetic anisotropy in dimethylthioformamide is the same as in dialkylamides, whereas it is inverted in the other thioamides. The effect of concentration has been studied with methyl N,N-dimethylthiocarbamate (I, R = CH,O), the only liquid compound in the series, in o-dichlorobenzene with the molar fraction of the solute varying from 1.0 to 0.200. The AF* values are almost independent of the concentration (Table I), whereas the AS* values show a gradual increase over a 5-eu range. The variation is not very much larger than the standard deviation and not too much weight can be attached to the trend. However, it is in the same direction as should be expected on the basis of the theory of Hatton and Richards,13 with an increasing proportion of the thioamide molecules forming complexes with the solvent. Correlation of Barriers with Calculated Bond Orders and n-Electrtm Energy Diflerences. It is naturally of considerable interest to select as the barrier for correlation with the molecular orbital quantities the thermodynamic parameter which is the best measure of the bonding energy in the molecule. According to Streitwieserl1* AF* is better in this respect than A H * . Hepler15has shown that under certain conditions AF measures rather well the internal enthalpy of activation. It is not possible to check if the conditions are fulfilled in the present systems, but the rather constant values of AF* in different solvents and concentrations show that this quantity is fairly independent of external influences and therefore it is chosen to represent the barriers. However, the differences between A F T and E , ( A H * ) are not very large and correlation with one of the latter parameters gives essentially the same result. As a consequence of the insensitivity of AF *, the well-known compensation of changes in A H and AS*16 is observed. For the amides, the AF*2ss.2 values of Rogers and Woodbrey6 are taken to represent the barriers. They are, with one exception, measured without solvent, but it is probable that they a t least fall in the same order as they would have done in o-dichlorobenzene. It is somewhat distressing that the barriers for dimethylformamide and dimethylacetamide differ by 3.6 kcal/mole, whereas the simple calculations predict the same barrier for them. However, as will be discussed later, the inductive effect Of the group is predicted to lower the barrier and hyperconjugation will work in the same direction. Therefore, dimethylformami& and dimethylthioformamide are chosen to represent the simple systems. From the data given
*
*
in ref 3, AF values of 26.6 kcal/mole (neat) and 24.0 kcal/mole (4001, in o-dichlorobenzene) can be calculated for dimethylthioformamide and the latter value can be used for correlation, even if the thioamide concentration is higher than in the present work. As a crude first approximation, a correlation waa attempted with the n-bond orders, as an increased C-N bond order should lead to a stiffening of the bond and an increase in the barrier. However, the correlation is not very good (Figure l), even if the bond order in all cases is higher in a thioamide than in the corresponding amide, which is also the case with the barriers, In particular, too low barriers are predicted for the compounds in which R contains double or triple bonds compared to those in whichR conjugates with lone pairs. n-Electron Energy Efect. In the transition state, the dimethylamino group is rotated 90’ out of the molecular plane around the C-N bond and therefore no 7r-electron interaction between the two parts of the molecule can take place. If this were the only effect on the energy of the molecule, the barrier should be measured by the loss in 7r-electron energy, AE,, following a 90’ rotation. M
r
=
E,,RCXNM~* - (E,,Rcx
+ 2aN)
(x= 0, s)
In this case, the effect of the substituent R on the barrier can be represented as the difference between
I
I
‘80 07!
20 -
lo -
1+ 06
-
04
-
*
The Journal of Ph&cal Chemistry
Figure 1. Correlation between AF* and ~ C N : 0, thioamide; +, amide; 1, R = H; 2, CHaO; 3, CH&; 4, C1; 5, C H 2 4 H ; 6, Ph; 7, N d ; 8, EtOCO; 9, AcNH.
(14) A. Btreitwieser, Jr., ‘‘Molecular Orbital Theory for Organic Chemists,” John Wiley and Sons, Inc., New York, N. Y.,1961, P 311 (15) L.G . Hepler, J . Am. Chem. SOC.,85, 3089 (1963). (16) J. E. Lemer, J . OW. Chem., 20, 1202 (1955); cf. also 0. Exner, Collection Czech. Chem. Commun., 29, 1094 (1964), and J. E. Leffler, J. fig.chem., 31,533 (1966).
BARRIERS TO INTERNAL ROTATION IN THIOAMIDES
its energies of interaction with the thiocarbonyl group and with the thioamide group. Since the interaction with the thioamids group involves cross-conjugation with the strongly conjugating amino group, this term can be expected to be less affected by R than the energy of interaction with the thiocarbonyl group. The result is that the barrier will be lowered more, the more strongly R can interact with the thiocarbonyl group. Thus, whereas tetramethylthiourea shows no line broadening due to exchange down to -60°, diminishing the electron-donating capacity of one amino group by substitution with an acetyl group leads to a barrier of 16.1 kcal/mole. For N,N-dimethy1-N'phenylthiourea (I, R = PhNH) a preliminary barrier of 11.5 kcal/mole (T, = 230'K) has been obtained. In general, a rather good correlation between AEr and AF* was found (Figure 2). For the thioamides a Ieast-squares calculation gave regression line AF = 67.9AE, - 19.9 with a correlation coefficient of 0.859. Model with Conjugating Heteroatom. The effect on ~ C and N AE, of increasing electron-donating capacity of a heteroatom substituent (Y in 11) can be clearly demonstrated by increasing ICCY with constant h y (Figure 3) or by decreasing h y with constant b y (Figure 4). In both cases X = S.
*
2323
20
15
t Figure 2. Correlation between AF* and a,. The regression line is calculated for the thioamides only. AE, is in unib of 8 ; for other symbols see Figure 1.
X
Y-c
// \
NR2
I1 In this connection it is of interest to compare the barriers for the thioamides with R = CH30 and R = CHSS. With the present parameters, both bond orders and AEr values predict the same barrier, whereas in fact for I, R = CHaS AF* is about 2 kcal/mole lower than for I, R = CHaO. This shows that alkylthio groups interact more strongly with the thiocarbonyl group than do alkoxy groups. The same order is valid for the carbonyl group, as is shown by the more qualitative observations reported by Valega.17 This difference can be ascribed to the greater polarizability of the substituent sulfur atom, which responds more strongly to the electron demand of the carbonyl and thiocarbonyl group, probably by contraction of the 3p orbitals. This points out a serious shortcoming of the simple HMO method and also of the w method. When the same parameters as have been employed here are used to calculate the interaction of RS and RO with phenyl, thiophene, and thiazole rings, the results predict stronger effects, e.g., electron displacements, with RS than with RO, whereas, as is well known, the reverse is
0.2 0.4 0.6 0.8 1.0 k CY
Figure 3. Plot of ~ C (-N ) and AE, ( - - - - ) V8. ~ C when Y hy = 1.5.
true, as is shown by the Hammett up values of -0.047 for CHaS and -0.268 for CH30.18 Obviously, the characteristics of the sulfur atom cannot even approximately be universally represented by only two parameters. The HMO method can only be expected to give reasonable agreement with observables within series of structurally rather closely related compounds. Thus, KramerlBhas shown that the barriers previously determined by himself and Gompperm for a series of (17) T.M.Valega, J . 070. Chem., 31, 1150 (1966). (18) H.H.J&B, Chem. Rev., 53, 191 (1953). (19) H.E.A. Kramer, Ann. Chem.,696, 28 (1966).
JAN SANDSTR~M
2324
AEn
Figure 4. Plot of p c (-~ ) and U T(- - - -) us. h y when lccy = 0.8.
N,N-dimethyl-P-acylenamines show a reasonable correlation with AET values from HMO calculations. The Inductive E f e c t . The influence of a purely inductive effect of Y has been studied by varying hc in 11, X = S. The A E , values (Table IV) show that a sub_____~
Table IV hC
PCN
AET
+0.2
0.452 0.455 0.459
0.651 0.636 0.622
0 -0.2
stituent with -I effect should increase the barrier and one with +I effect lower it. The former effect could, at least partly, explain the relation between the barriers of dimethylformamide (21.0 kcal/mole), dimethylacetamide (17.4 kcal/mole), and dimethylpmpionamide (16.7 kcal/mole), and the latter effect the relation between the barriers of dimethylacetamide and dimethyltrifluoroacetamide (17.6 kcal/mole) . For the latter molecule, models indicate a small steric hindrance to the planar state, which, as discussed later, should lower the barrier, and this effect should be still greater in dimethyltrichloroacetamide (14.9 kcal/mole) . All barriers are from ref 6. Steric Effects. It is obvious that the barriers to internal rotation must be considerably sensitive to steric effects. Crowding in the planar state will raise its energy and thus lower the barrier, and Mannschreck21 and Staab22have shown that crowding in the transition state can raise its energy to such an extent that The J o u r 4 of Physical C h i s t r y
individual rotamers can be isolated. Dahlqvist and Fors6n2ahave shown that the AF values for benzaldehyde, p-methoxybenzaldehyde, and p-dimethylaminobenzaldehyde increase in the same order as the AET values obtained by the HMO method, but that the ratio between the highest and the lowest AET value is only 1.05 whereas the ratio between the corresponding barriers is 1.37. When 2-furanaldehyde was compared with the benzaldehydes, no good correlation was obtained. The authors ascribe this shortcoming to important contributions to the barriers from nonbonding interaction. In the thioamide series the ratio between the highest and the lowest AE= value is 1.24 and between the corresponding AF* values 1.49. Thus, the relation is somewhat better than for the benzaldehydes, though the calculations still underestimate the differences between the barriers of the different compounds. On the other hand, insertion in A E r of acceptable P values of the order of 30 kcal/mole gives AF* values ranging from 15 to 19 kcal/mole, which means that ?r-electron effects can account for the larger part of the barriers. There are also good reasons to believe that the dimethylamino group is comparatively free from interactions with neighboring groups in most of the cases investigated here. It has been foundz4in a series of N-methylthioamides (111, R = H, CH3,C2H5,i-C3H,, and t-C4H,) that the rotamer (IIIa) predominates strongly in solution. This ob-
*
PCN
S R-C
// \
S
R-C N-CH,
// \
N-H
/
H
IIIa
IIIb
servation shows that no repulsion of importance can exist between the sulfur atom and the methyl group. A similar observation has been made for some oxygen analogs.26 Molecular models and known van der Waals radii show that no interaction needs to take place between the other methyl group and R when R is H, CHI, CH30, CHIS, and C1. On the other hand, it has been shown that when R = ( C H ~ ) Z NPh,27 ,~~ (20) H.E. A. Kramer and R. Gompper, 2.Physik. Chem. (Frankfurt), 43, 292 (1964). (21) A. Mannschreck, Tetrahedron Letters, 1341 (1965). (22) H.A. Staab and D. Lauer, ibid., 4593 (1966). (23) K-I. Dahlqvbt and S.Forsh, J. Phys. Chem., 69, 4062 (1965). (24) J. Sandstrom and B. Uppstrom, submitted for publication. (25) L. A. LaPlanche and M. T. Rogers, J. A m . Chem. SOC.,86, 371 (1964).
BARRIERS TO INTERNAL ROTATION IN THIOAMIDES
or EtOC0,28considerable steric effects are at work, which in the two latter molecules probably lead to rotation of the planar or nearly planar thioamide group out of the molecular plane. It is still possible, by a simultaneous rotation around C-N and C-C bonds, to transfer a methyl group between the two nonequivalent sites without having to press it past the obstructing atom. However, in this case the shape of the potential surface will probably be different from that when the dimethylamino group is unhindered, but the moderate agreement between calculated and experimental barriers indicates that this effect is not of overwhelming importance. Calculation of Barriers in Amidinium Salts, Amidines, and Enamines. An extrapolation from the relation between amides and thioamides leads to the assumption that mesomeric interaction between an amino group and a C=X group will decrease with decreasing electronegativity of X and with increasing tendency of C and X to form a double bond. The effect of the nature of X can be demonstrated with the same model as before (11). With Y = alkyl, calculations have been performed with parameters for X corresponding + to NR2 (amidinium ion, h~ = 1.5), NR (amidine), and CH2 (enamine). The results (Table V) show a crude correlation between experimental and calculated quantities. For aniline the same parameters predict almost the same barrier ( ~ C N= 0.314, AE, = 0.396p) as for enamines. Evans29 has found the barrier in aniline to be 3.54 kcal/mole and a value of the same
2325
Table V X
PCN
AE,
-AF*
S
0.455
0
0.422
0.636 0.532
24.50 19.20
AFtl NR CHz
0.484 0.363 0.316
0.632 0.449 0.394
19.80
Calculated from data in ref 4. No value known. above -40".
* No
b C
exchange broadening
order of magnitude should be expected for the enamines. It is true that the nonplanarity of the aniline30and probably also of the enamine molecule should require a higher h~ and a lower ~CCN value, but the resultant lowerings of the barriers are parallel for the two systems and the above arguments are still approximately valid.
Acknowledgments. The author is grateful to The Swedish Natural Sciences Research Council for financial support and to Fil. Mag. Kerstin Malmqvist for programming the nmr calculations. (26) M. J. Jansaen, Rec. Tmv. C h a . , 79, 454 (1960). (27) J. Sandstrom, Acta Chem. Smnd., 16, 1616 (1962); J. Sandstrom and B. Uppstrom, ibid., 19, 2432 (1965). (28) B. Persson and J. Sandstrom, ibid., 18, 1059 (1964). (29) J. C. Evans, Spectrochim. Acta, 16, 428 (1960). (30) D.G. Liater and J. K. Tyler, C h a . Commun., 3, 152 (1966).
Volume 71, Number 7 June 1967
