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The infra-red spectra and hydrogen bonding of heterocyclic thioamides
Spectmchhica Acta, Vol. 288, pp.1899to 1907. Pergamon Prem 1972. Printed in Northern Ireland
The ix&a-red specba and hydrogen bonding thi0tiat2s
of heterocyclic
P. J. F. GRIFFITHS and G. D. MORGAN Depsrtment of Chemistry, University of Wales Institute of Science and Technology, Cardiff CF13NTJ, Wales
and BRYAN ELLIS Department of Glass Technology, University of Sheffield, SheffieldSIO 2TZ, England (Received20 Janwavy 1972) Al&&-Hydrogen bonding of the heterocyclia thioamides thiazolidine-2-thione(I), benxthiazoline-2-thione(I1)and 4-methylthiaxoline-2&ione(III) is interpretedin terms of the formation of cyclio dimers, which favours co-operative proton transfer and resonance stabilixation. . . . The dmermatlon constants, &,,, have been determined for each of the three compounds in solution in carbon tetrachloride at 26 f l”C, [(I): /?e,, = 88 f 9 M-r; (II): /lse = 670 rt 60 M-1, (III): /la0 = 6540 & 600 M-l]. The strengths of the hydrogen bonds, which have been e&meted from Av = (~xxr,,~ - vxx...s), increaseintheorder < II < III. Thestrength of the hydrogen bond of III is approximately equal to that of thiopyridone.
THE STUDYof the hydrogen bonding of thioamides is important because of the possibility of co-operative proton transfer if cyclic dimers are formed. Co-operative proton transfer was discussed many years ago by HUWXNS[l] who used the term “synchronized oscillation resonance”. This effect is important in the stabilization of the hydrogen bonds of the base pairs in nucleic acids, of which the pair adenine and thymine ia an example. The present work is concerned with hydrogen bonding to sulphur. The i.r. spectra of thiazolidine-2thione(I), benzthiazoline-2-thione(II), and 4methylthiazoline-2-thione(II1) in solution in carbon tetrachloride were reported by FLETT[2].
In these spectra there was a sharp band near 3420 cm-l, which was assigned to the free N-H of the monomer and strong bands between 2900 and 3100 cm-l, which [l] M. L. HUQGINS, Aragew. Chm. Intern.Ed. Engl. [2] M. ST. C.F'LETC, J. Chem. BOG.847 (1953).
1899
10,(3), 147 (1971).
1900
P. J. F. GRIFFITES, G. D. Moso~ and B. ELLIS
were assigned to the hydrogen bonded N-H present in complexes, such as the dimer, which for II is represented by IV, in equilibrium with the monomer.
(IV) For II in carbon tetrachloride solution the relative intensities of the free N-H band (3420 cm-l) and the hydrogen bonded N-H bands (2900-3100 cm-i) varied with concentration. ELLISand GRIFFITHS [3] discussed the strength of the hydrogen bonding in thioamides as revealed by the effect of solvent on the U.V.spectra of these compounds. The i.r. spectrum of the hydrogen bonded species was reported in greater detail for II by BEJX,AMYand RO~ASCH[4], who stated that the spectrum for the solid was essentially the same as that for a solution in carbon tetrachloride, indicating that the hydrogen bonded species had the same structure in both phases. BELLAMYand ROQASCH [4] discussed their results for II in relation to those for other molecules in which co-operative proton transfer could occur. For such an effect to be important it would seem to be essential for dimers to be formed. The formation of dimers was presumed by BELLAMY and RO~ASCH[a] from a report of a space group determination by JEFFREY [5] in which it was tentatively suggested that the molecules of II were arranged in pairs in the unit cell. Contradictory evidence was produced by TASHPULATOV, ZVONI(OVA and ZHDANOV [6], who made a threedimensional crystal structure determination of II, based on 672 reflections using Gu radiation, and concluded that the molecules of II were arranged in helical chains about a two-fold screw axis with an exceptionally short intermolecular sulphurnitrogen distance and hence unusually strong hydrogen bonds. Recently the crystal structure of II, based on 1536 reflections using MO Ra radiation, has been redetermined by CHESICKand DONOHUE[7] who show that the molecules form centro-symmetric hydrogen bonded dimers with no unusually short hydrogen bonds. BELLAMYand ROQASCH[4] pointed out ten years ago that stability constants were not available for II and the little work done on thiomides was referred to by KULEVSKYand I?ROEHLICH [8]. The association of II was iuvestigated by HOPKINS and HUNTER[9] as part of a general study of mesohydric tautomerism. However, the association factors, a = JZ/(Formula Weight) where M is the molecular weight [3] BRXAN EIAIS and P. J. F. GRIFFITHI,Spectrochim. Acti a, 2006 (1966). [4] L. J. BELLAMY and P. E. ROQASCH, Proc. Roy. Sot. (London) A957,98 (1960). [6] G. A. JEFFREY. In H. P. KOCH, J. Chem. Sot. 401 (1949). [0] Iu. TAB EPTJLATOV,z. v. ZVONEOVA md G. 8. ZHDANOV, Soviet Phys. c%y&. a,33 [7] J. P. CHESICK md J. DONOFIUE, Acta Or@. B27, 1441 (1971). [S] N. KULEVSKY and P. M. F’ROEHLICH, J. Am. C&m. Sot. 89,4839 (1967). [9] G. HOPIXINS and L. HUNTER, J. Chem. Sot. 638 (1942).
(1957).
The infra-red spectra and hydrogen bonding of hetmocyclic thioamidee
1901
could not be explained quantitatively. From these observations and our own U.V. spectroscopic studies it ww clear thaf further work in this area was desirable. The objectives of fhe present work were fo show tha;t dimers are present in solution and to determine the self-association constants for I, II and III in carbon tetrachloride solution. The strengths of the hydrogen bonds were estimated from the spectral shifts which occur when hydrogen bonds are formed.
in solution
in nrcphthalene,
EXPEIUMENTAL COvn~?& Thiazolidine-2-thione(I), m.p. 104”, (lit. 107” [lo]) donatid by the Anchor Chemical Co. Ltd., benzthiazoline-2-thione(II), m.p. 178-179’ (lit. 177-179’ [ll]) donated by Monsanto Chemicals Ltd., and 4-methylthiazoline-2-thione(III), m.p. 87” (lit. 90” [12]) donated by I.C.I. Ltd., were used after recrystallization to constant melting point. The solutions of the thioamides were prepared by semimicro weighing. An ulfrasonic bath was used to obtain the solutions because of the low solubility of the compounds. The solutions were checked carefully prior to measurement to ensure that no precipitation had occurred. Concentrations were at intervals of O-1 x lo-* M in the range O-5-1 *O x 1O-s M. Studio at concentrations much above lOa M could not, be made because of the low solubility of the compounds ( <5 x 10-s M). The spectroscopic grade carbon tetrachloride used was stored over a 4 A molecular sieve and distilled prior to use, to remove water which is normally present to the extent, of NO-01o/o unless the solvent has been specially dried [13]. By comparison of the i.r. spectrum (3400-4000 cm-l) of the dried solvent with that of undried spectroscopic carbon tefrachloride, using 10 cm cells, the’ amount of water was found to be less than O-0014% (~10~ M) [13, 141. The residual water in the carbon tetrachloride was af a concentration of ~10-8 M, so that there was a possibility of water-thioamide interaction. CEEHSTUNet al. [ 151have shown that water is not significanfly self-associated in carbon tetrachloride over a range of temperatures (15-45’C) and concentrations (O-15 x 1O-8M). In the present work if was found that when the carbon tetrachloride was nearly saturated with water (~0-01~/~) the relative intensities and frequencies of the water fundamentals, at 3620 and 3705 cm-l were the same whether measured for water in carbon fetrachloride alone or in the presence in solution of any one of the thioamidea, at concentrations of 0.5 x lOa M and O-25 x 10-J M. There was [IO] E. H. RODD (Editor), Chmtitry of
Carbon Coqwounda, Vol. IV, Pt. A, p. 408. Elmvier, New York (1967). [l l] I. HEILBRON and H. M. BUNBURY, Dictionmy of Organic Compounde, Eyre and Spottiswood, London (1953). [12] E. H. RODD (Editor), Chemaetry of Carbon C%nnpoud, Vol. IV, Pt. A, p. 396. Else&r,
New York (1957). [13] W. ESS~BORN and B. HAMPEL, P.S.U. BuUetin No. 17 (September 1967). [Ia] W. ESSELBORNand B. HABIPEL,Private communication. [ 161 S. D. CHRISTLAN,H. E. AFXSPBUN~and J. R. JOHNSON,J. Chem. Sot. 1896 (1983).
1902
P. J. F. GRIFFITHS,G. D. MORUANand B. ELLIS
no sign&ant change in the apparent molar extinction coefficient of the free N-H band of I, II and III when the spectra were measured at four concentrations of water in the range 0~070-O~OO14°/o. From these results it was concluded that water-solute interaction was negligible. Spectra The measurements of the free N-H band and the water fundamental vibrations were made on an SP700 spectrophotometer. 10 cm cells were used to examine the effect of water in the carbon tetrachloride and 1 cm cells to study the self-association of the thioamides. All measurements were made at 25 f 1°C. Wave numbers are given to f 10 cm-l. Measurements of the hydrogen bonded region were made on the solids at 1 o/oin KBr discs and solutions in Infrasil cells on a Per&in-Elmer 521 spectrophotometer. RESULTS AND Discussion
Infrared spectra The general features of the i.r. spectra of the thioamides I, II and III in solution in carbon tetrachloride are a readily observable sharp band near 3410 cm-l, with a half-band width of 20 cm-l, due to free N-H and a hydrogen bonded region extending over the range 2700-3100 cm-l. The broad hydrogen bonded region has resolvable overlapping sub-bands. The general characteristics of such bands were discussed by SHEPPARD [16] and in detail for II by BELLAMY and ROC+ASCH[4] in terms of Fermi resonance. For II there is little detectable difference between the hydrogen bonded regions of the spectra in carbon tetrachloride and as a solid (in a KBr disc), in agreement with the observation of BELLAMY and ROGANH [a]. This is primary evidence for the existence of II as cyclic dimers (IV) in solution. If the cyclic dimers opened to form linear chain oligomers in addition to monomer, some change in the spectra in this region would be expected. For pyrazole the change from hydrogen bonded helical chains in the solid [17] to monomer, cyclic dimer and trimer in solution in carbon tetrachloride is accompanied by an increase of about 30 cm-l in the positions of the maxima in the hydrogen bonded region [18]. It is di&ult to be certain that there are no changes in the i.r. spectra of I and III in passing from the solid to solution in carbon tetrachloride, because of the presence of C-H vibrations in the 2940-2840 om-1 region which overlap the hydrogen bonded N-H absorption. However, there do not appear to be any major differences between the solid and solution spectra. (We have been unable to trace crystal structure determinations for either I or III.) In Table 1 the positions of the free N-H and the multiple overlapping sub-bands are given for I, II and III for both solid dispersed in a KBr disc and solution. In general there is agreement to within ~15 cm-l between the present results for II and those of BELLA~ and ROGASCH[4]. Further detailed discussion of the hydrogen bonded absorption is complicated because of the multiplet structure but the band-widths are comparable (I N 650 [la] N. SHEPPARD,Hydrogen Bonding, p. 85. (Editedby D. HADZI) Pergamon,London (1959). [17] H. W. EHRLIOH,Acta Cry&. 18, 946 (1960). [18] D. M. W. ANDERSON,J. L. DUNCAN and F. J. C. ROSSOTTI,J. Chm. Sot. 140(1961).
The infrrt-redspectra and hydrogen bonding of heterocyclicthioamides
1903
Table 1. N-H absorption bands (cm-l) of thiazoline-2-thione(I), benzthiazoline-2-thione(I1) and 4.methylthiwoline-2thione(III) in solid (KBr disc) and solution in carbon tetraohloride II
I Solid
3130 3000 2930 2866 2846 2700
Solution 3420 3130 298O(sh) 2930 2870 2860 -
Solid 3190 3130(sh) 3100 3070 3030 3010(sh) 2960 2880 2840 2760 2730 2060
III Solution 3410 3190(sh) 3110 3070 3036 302O(sh) 2960 2890 2840 2760 2730 2680
Solid
3186 3116(sh) 3086 3036 3000 296O(sh) 2916 2900 2866 2836 2710 266O(sh)
Solution 3410 3186(sh) 3120 3086 3040 3006 2960 2930(sh) 2906 2876 2860(s) 2710 266O(sh)
sh denotes shoulder.
II h) 760 cm-l, III N 800 cm-l). The band intensity of III is about 60% tha.n that of II which is considerably larger than that of I. The approximate band centres axe: I, 3020; II, 2960; III, 2880 cm-l. There have been several attempts to correlake Av(v~ - v~.,.~) with the energy of a hydrogen bond, and it is generally reoognised that such correlations should be used cautiously [19-241. However, for the compounds I, II and III both band intensities and Av indicate that the strength of the hydrogen bonding increases in the order I < II < III. This is in agreement with the dimerization constants (see later) and the conoluaion of ELLISand GRIFFITHS [3] from a study of the effect of hydrogen bonding on the U.V. spectra of these compounds. The hydrogen bonded band centre for III does not difFersignificantly from the band centre for thiopyridone [4] indicating that the N-H . . . S hydrogen bond in III is aa strong as that in thiopyridone, one of the strongest so far recorded [25,26]. cm-l,
greater
Determhation of association constants As 8 consequence of the low solubility of the thioamides I, II and III the choice of solvent for the study of association constants was restricted. Carbon tetrachloride was found to be the most suitable. The ~88 of carbon tetrachloride for such studies has been the subject of much discussion because of the possibility of hydrogen R. M. BADGER and 8. H. BAUJSR,J. Chem. Phyu. 6, 839 (1937). R. M. BADGER,J. Chm. Phys. 8,288 (1940). N. D. COUGESIXALZ, J. Chern.Phyu. 18,978 (1960). E. R. LIFPINCOTJIand R. SCHILOEDEFG, J. Clrem.Phyu. 28,lOQQ (1966). and mcdexdarStructure,p. 413. (Edited by MANSEL H. E. HALUX, Irafra-redSpectrosco2ry DAVIES) Elsevier, London (1963). Chapter 8, Methuen, London (1968). [24] L. J. BELLABSY,Advanca in Infra-red .9pect~o.9copy, [26] B. R. -OLD, Acta CY’y8t. 6, 707 (1963). [26] N. KULEV~KYand W. REINEKE, J. Phys. Chem. 73,3339 (1968). [lQ] [20] [21] [22] [23]
1904
P. J. F.
GRIFFITHS, G.
D. MORGAN and B. Ems
bonding to the chlorine atoms [27-311.In the present study we conclude that hydrogen bonding beween the monomeric form of the thioamides and carbon tetrachloride does not occur because water, a stronger hydrogen-bonding base, does not affect the i.r. spectra [see EXPERIMENTAL].The degree of solvation of the monomer and dimer would not be expected to be the same of course. From the similarity of the i.r. spectra of II in both the solid state and carbon tetrachloride solution it is concluded that the dimeric form is predominant in solution. If only the monomer-cyclic dimer equilibrium 2MBT
= (MBT),
(1)
is involved the stability constant &,[32] is defined by B 20
=
IWBThl [MBT12
and the total concentration, B, of thioamide is given by B = b + 2/9,,b2
(3)
where b is the concentration of monomer [32]. The concentration of monomer was determined by measurement of the apparent molar absorptivity of the free N-H band at ~3410 cm-l. If Beer’s law is obeyed b = E . B/E~ where .a0is the molar absorptivity of the free N-H band and E is the apparent molar absorptivity measured at a stoiohiometric concentration B. Direct determination of a0is not possible because there is still significant hydrogen bonding even at the lowest concentration at which the free N-H absorbance could be determined accurately. Substitution for b in equation (3) and rearrangement of terms yields
so that pa0 may be calculated if a plot of l/e against EB is linear. The l/e against EB plots for I, II and III are linear (Fig. 1). The slopes and intercepts were calculated by the method of least squares and the values of a0 and pzo calculated from these are given in Table 2. This treatment has been used previously by HARRISand HOBBS[33] and VINO(XRADOV [34]. In general, if cyclic dimers and linear oligomers of variable size are formed then a complicated relationship is obtained B = b + 2p2,b2 + IZ npnObn
(6)
[27] V. LIDDEL andE.D.B~~mm,Spectrochim.Acta lo,70(1967). [28] P. HWSKENS, R. HENRY and G. GILLWOT, Bull. Sot. Chtkn. France 720 (1962). [29] D. A. IBBITSON and L. F. MOORE, J. Chem. Sot. (B) 76 (1967). [30] A. BALL md J. L. WOOD, Spectmchim. Acta 24& 1108 (1968). [31] A. N. ~ETOHER, J. .Wqp. Chem. 73,2217(1969). [32] F. J. C. ROSSOTTI and H. ROSSOTTI, The determination of stability cm&an& McGraw-Hill, New York (1961). [33] J. T. HARRIS JR. and M. E. HOBBS, J. Am. Chms. Sot. 76, 1419 (1964).
[34] S. N. VINOURADOV, Can. J. Chews. 41, 2719 (1963).
The i&cl-red spectra and hydrogen bonding of heterooyolicthioamides
41 30
I
50
I
I
70
90
1
I
I
I
I
110
130
150
170
190
tEl(xl03)
Fig. 1. (a) Plot of l/e vs. EB accordingto equation (4).
12 -
II -
IO -
6-
?-
65
/ I 30
34
I 36
I 42
I
I
I
I
I
I
46
50
54
56
62
66
(I /a
x
10-3
Fig. 1. (b) Plot of data for I according to the rearrangedequation (4): 1
60
EB
esB
-=---
%O 60
because of the small slope for these date in (a). (Slope = e. = 192)
1905
1906
P. J. F. GaIFPITHS,G. D. MORUANand B. ELLIS Table 2. Molar absorptivities (Q) and self association constants (j&) Compound
80
I II III
192 f 10 196 & 10 200 f 10
88 f 9 670 f 60 5640 f 600
The values of so and Be0in this table were obtained by a least squares numerical fit to equation (4) and not by a graphical method.
Equation (6) is reduced to an expression similar to equation (4) if cyclic and linear dimers only sre in equilibrium with monomer [35]. Thus, when a linear plot of l/e sg&st EB ia obtained it is not established unequivocally that cyclic dimer only is present. However, from the similkty of the i.r. absorption in the hydrogen bonded region for II both in the solid and solution it appears to be definitely established that cyclic dimers are present in solution. This is an alteration of our previous point of view [3], based on the crystal structure proposed by ZVONKOVA et al. [6], that linear oligomers were present in solution. While the evidence is not so dekite it appears that I and III exist predominantly as cyclic dimers also in solution. The small amounts of open-chain dimer which might possibly be present would also absorb at the monomer frequency. It is unlikely, however, that the correction to ,&, would be greater than about 10% [36]. From the geometry of the molecules and the crystal structure determination [7] for II it is seen that the cyclic dimers have linear N-H . . . S hydrogen bonds and we conclude thet essentially, in carbon tetrachloride solution (26’C, 0.5-l x 10-z M) the equilibrium is between cyclic dimer and monomer. Increase in temperature results in sn inoresse in intensity of the 3410 cm-i bsnd due to free N-H rend&reduction in that of the broad hydrogen bonded N-H band, for each of the three compounds. This indicates that the enthalpy changes on dimerization are negative, if it is assumed that the temperature coefficient of e,, is small [37]. In each of the three compounds, although the hydrogen bonding involves the S II -NH-Cgroup of atoms, the effect of the remsinder of the moleoule is quite pronounced. Steric effects csnnot be involved snd the hydrogen-donor capacity is governed largely by the extent of conjugstion of the heterocyclic ring with the S -NH-(!!group. For thiazolidine-2-thione there is some conjugation with the S atom in the heterocycle lectding to a. slight enhancement of hydrogen bonding [36] R. C. LORDand T. J. Poaao, 2. EZektrodwm. 64,672 (1960). [36] N. KULEVSKY and P. M. F~OEEIJCH, J. Am. Chem. Sot. S&4839 (1967). 1371 G. C. Pnawr~~ and A. L. MCCLZELLILN, TlaeHyhogen Bond, p. 76. W. H. Freeman, San Fran&co and London (1960).
The i&k-red spectra and hydrogen bonding of hetsrocyclic thioamides
1907
capacity. For 4-methylthiazoline-2-thione, however, there is much more extensive conjugation with the double-bond between C4 and C5, with consequent considerable increase in hydrogen bonding capacity. The feebly electron donating methyl group would not be expected to have a significant effect. With benzthizoline-2-thione the S II
extent of conjugation of the double bond between C4 and CS with the -NH-Cgroup is reduced by the conjugation in the six membered ring and there is a concomitant reduction in hydrogen bonding’capacity. This demonstrates the considerable importance of the covalent contribution to the hydrogen bond in these compounds. The N or S methyl derivatives of these compounds can be prepared easily, indicating that chemically they exhibit tautomeric behaviour although normslly they exist predominantly in the thione form [Z, 31. In view of the dual nature of the chemical reactivity of these compounds it is probable that the difference between the energies of the thione and thiol forms is not large so that co-operative proton transfer [38] in the excited state, between the forms V and VI
‘N
I ?
A‘S
i
Q
would be expected to occur to a larger extent than customarily with NH . . . S bonds. In accordance with this view is the fact that the hydrogen bonded band widths are similar, although the positions of the band centre and pa0 values are consistent with the increasing strength of hydrogen bond formed in going from I to III, indicating that the lttrge band-width is not merely a consequence of the usual relationship between AY and vt [39]. Acknowledgements-We are grateful to Messrs. E. WANNELLand J. CANNARDfor experimenti &k3t&IC0. [38] C. G. CANNON,S~ectrochim.Acta 10,341 (1968). [39] G. C. PIXENTELand A. L. MCCLELLAN,2% Hydrogen Bond, Francisco ad London (1960).
p. 94.W. H. Freeman, San
The ix&a-red specba and hydrogen bonding thi0tiat2s
of heterocyclic
P. J. F. GRIFFITHS and G. D. MORGAN Depsrtment of Chemistry, University of Wales Institute of Science and Technology, Cardiff CF13NTJ, Wales
and BRYAN ELLIS Department of Glass Technology, University of Sheffield, SheffieldSIO 2TZ, England (Received20 Janwavy 1972) Al&&-Hydrogen bonding of the heterocyclia thioamides thiazolidine-2-thione(I), benxthiazoline-2-thione(I1)and 4-methylthiaxoline-2&ione(III) is interpretedin terms of the formation of cyclio dimers, which favours co-operative proton transfer and resonance stabilixation. . . . The dmermatlon constants, &,,, have been determined for each of the three compounds in solution in carbon tetrachloride at 26 f l”C, [(I): /?e,, = 88 f 9 M-r; (II): /lse = 670 rt 60 M-1, (III): /la0 = 6540 & 600 M-l]. The strengths of the hydrogen bonds, which have been e&meted from Av = (~xxr,,~ - vxx...s), increaseintheorder < II < III. Thestrength of the hydrogen bond of III is approximately equal to that of thiopyridone.
THE STUDYof the hydrogen bonding of thioamides is important because of the possibility of co-operative proton transfer if cyclic dimers are formed. Co-operative proton transfer was discussed many years ago by HUWXNS[l] who used the term “synchronized oscillation resonance”. This effect is important in the stabilization of the hydrogen bonds of the base pairs in nucleic acids, of which the pair adenine and thymine ia an example. The present work is concerned with hydrogen bonding to sulphur. The i.r. spectra of thiazolidine-2thione(I), benzthiazoline-2-thione(II), and 4methylthiazoline-2-thione(II1) in solution in carbon tetrachloride were reported by FLETT[2].
In these spectra there was a sharp band near 3420 cm-l, which was assigned to the free N-H of the monomer and strong bands between 2900 and 3100 cm-l, which [l] M. L. HUQGINS, Aragew. Chm. Intern.Ed. Engl. [2] M. ST. C.F'LETC, J. Chem. BOG.847 (1953).
1899
10,(3), 147 (1971).
1900
P. J. F. GRIFFITES, G. D. Moso~ and B. ELLIS
were assigned to the hydrogen bonded N-H present in complexes, such as the dimer, which for II is represented by IV, in equilibrium with the monomer.
(IV) For II in carbon tetrachloride solution the relative intensities of the free N-H band (3420 cm-l) and the hydrogen bonded N-H bands (2900-3100 cm-i) varied with concentration. ELLISand GRIFFITHS [3] discussed the strength of the hydrogen bonding in thioamides as revealed by the effect of solvent on the U.V.spectra of these compounds. The i.r. spectrum of the hydrogen bonded species was reported in greater detail for II by BEJX,AMYand RO~ASCH[4], who stated that the spectrum for the solid was essentially the same as that for a solution in carbon tetrachloride, indicating that the hydrogen bonded species had the same structure in both phases. BELLAMYand ROQASCH [4] discussed their results for II in relation to those for other molecules in which co-operative proton transfer could occur. For such an effect to be important it would seem to be essential for dimers to be formed. The formation of dimers was presumed by BELLAMY and RO~ASCH[a] from a report of a space group determination by JEFFREY [5] in which it was tentatively suggested that the molecules of II were arranged in pairs in the unit cell. Contradictory evidence was produced by TASHPULATOV, ZVONI(OVA and ZHDANOV [6], who made a threedimensional crystal structure determination of II, based on 672 reflections using Gu radiation, and concluded that the molecules of II were arranged in helical chains about a two-fold screw axis with an exceptionally short intermolecular sulphurnitrogen distance and hence unusually strong hydrogen bonds. Recently the crystal structure of II, based on 1536 reflections using MO Ra radiation, has been redetermined by CHESICKand DONOHUE[7] who show that the molecules form centro-symmetric hydrogen bonded dimers with no unusually short hydrogen bonds. BELLAMYand ROQASCH[4] pointed out ten years ago that stability constants were not available for II and the little work done on thiomides was referred to by KULEVSKYand I?ROEHLICH [8]. The association of II was iuvestigated by HOPKINS and HUNTER[9] as part of a general study of mesohydric tautomerism. However, the association factors, a = JZ/(Formula Weight) where M is the molecular weight [3] BRXAN EIAIS and P. J. F. GRIFFITHI,Spectrochim. Acti a, 2006 (1966). [4] L. J. BELLAMY and P. E. ROQASCH, Proc. Roy. Sot. (London) A957,98 (1960). [6] G. A. JEFFREY. In H. P. KOCH, J. Chem. Sot. 401 (1949). [0] Iu. TAB EPTJLATOV,z. v. ZVONEOVA md G. 8. ZHDANOV, Soviet Phys. c%y&. a,33 [7] J. P. CHESICK md J. DONOFIUE, Acta Or@. B27, 1441 (1971). [S] N. KULEVSKY and P. M. F’ROEHLICH, J. Am. C&m. Sot. 89,4839 (1967). [9] G. HOPIXINS and L. HUNTER, J. Chem. Sot. 638 (1942).
(1957).
The infra-red spectra and hydrogen bonding of hetmocyclic thioamidee
1901
could not be explained quantitatively. From these observations and our own U.V. spectroscopic studies it ww clear thaf further work in this area was desirable. The objectives of fhe present work were fo show tha;t dimers are present in solution and to determine the self-association constants for I, II and III in carbon tetrachloride solution. The strengths of the hydrogen bonds were estimated from the spectral shifts which occur when hydrogen bonds are formed.
in solution
in nrcphthalene,
EXPEIUMENTAL COvn~?& Thiazolidine-2-thione(I), m.p. 104”, (lit. 107” [lo]) donatid by the Anchor Chemical Co. Ltd., benzthiazoline-2-thione(II), m.p. 178-179’ (lit. 177-179’ [ll]) donated by Monsanto Chemicals Ltd., and 4-methylthiazoline-2-thione(III), m.p. 87” (lit. 90” [12]) donated by I.C.I. Ltd., were used after recrystallization to constant melting point. The solutions of the thioamides were prepared by semimicro weighing. An ulfrasonic bath was used to obtain the solutions because of the low solubility of the compounds. The solutions were checked carefully prior to measurement to ensure that no precipitation had occurred. Concentrations were at intervals of O-1 x lo-* M in the range O-5-1 *O x 1O-s M. Studio at concentrations much above lOa M could not, be made because of the low solubility of the compounds ( <5 x 10-s M). The spectroscopic grade carbon tetrachloride used was stored over a 4 A molecular sieve and distilled prior to use, to remove water which is normally present to the extent, of NO-01o/o unless the solvent has been specially dried [13]. By comparison of the i.r. spectrum (3400-4000 cm-l) of the dried solvent with that of undried spectroscopic carbon tefrachloride, using 10 cm cells, the’ amount of water was found to be less than O-0014% (~10~ M) [13, 141. The residual water in the carbon tetrachloride was af a concentration of ~10-8 M, so that there was a possibility of water-thioamide interaction. CEEHSTUNet al. [ 151have shown that water is not significanfly self-associated in carbon tetrachloride over a range of temperatures (15-45’C) and concentrations (O-15 x 1O-8M). In the present work if was found that when the carbon tetrachloride was nearly saturated with water (~0-01~/~) the relative intensities and frequencies of the water fundamentals, at 3620 and 3705 cm-l were the same whether measured for water in carbon fetrachloride alone or in the presence in solution of any one of the thioamidea, at concentrations of 0.5 x lOa M and O-25 x 10-J M. There was [IO] E. H. RODD (Editor), Chmtitry of
Carbon Coqwounda, Vol. IV, Pt. A, p. 408. Elmvier, New York (1967). [l l] I. HEILBRON and H. M. BUNBURY, Dictionmy of Organic Compounde, Eyre and Spottiswood, London (1953). [12] E. H. RODD (Editor), Chemaetry of Carbon C%nnpoud, Vol. IV, Pt. A, p. 396. Else&r,
New York (1957). [13] W. ESS~BORN and B. HAMPEL, P.S.U. BuUetin No. 17 (September 1967). [Ia] W. ESSELBORNand B. HABIPEL,Private communication. [ 161 S. D. CHRISTLAN,H. E. AFXSPBUN~and J. R. JOHNSON,J. Chem. Sot. 1896 (1983).
1902
P. J. F. GRIFFITHS,G. D. MORUANand B. ELLIS
no sign&ant change in the apparent molar extinction coefficient of the free N-H band of I, II and III when the spectra were measured at four concentrations of water in the range 0~070-O~OO14°/o. From these results it was concluded that water-solute interaction was negligible. Spectra The measurements of the free N-H band and the water fundamental vibrations were made on an SP700 spectrophotometer. 10 cm cells were used to examine the effect of water in the carbon tetrachloride and 1 cm cells to study the self-association of the thioamides. All measurements were made at 25 f 1°C. Wave numbers are given to f 10 cm-l. Measurements of the hydrogen bonded region were made on the solids at 1 o/oin KBr discs and solutions in Infrasil cells on a Per&in-Elmer 521 spectrophotometer. RESULTS AND Discussion
Infrared spectra The general features of the i.r. spectra of the thioamides I, II and III in solution in carbon tetrachloride are a readily observable sharp band near 3410 cm-l, with a half-band width of 20 cm-l, due to free N-H and a hydrogen bonded region extending over the range 2700-3100 cm-l. The broad hydrogen bonded region has resolvable overlapping sub-bands. The general characteristics of such bands were discussed by SHEPPARD [16] and in detail for II by BELLAMY and ROC+ASCH[4] in terms of Fermi resonance. For II there is little detectable difference between the hydrogen bonded regions of the spectra in carbon tetrachloride and as a solid (in a KBr disc), in agreement with the observation of BELLAMY and ROGANH [a]. This is primary evidence for the existence of II as cyclic dimers (IV) in solution. If the cyclic dimers opened to form linear chain oligomers in addition to monomer, some change in the spectra in this region would be expected. For pyrazole the change from hydrogen bonded helical chains in the solid [17] to monomer, cyclic dimer and trimer in solution in carbon tetrachloride is accompanied by an increase of about 30 cm-l in the positions of the maxima in the hydrogen bonded region [18]. It is di&ult to be certain that there are no changes in the i.r. spectra of I and III in passing from the solid to solution in carbon tetrachloride, because of the presence of C-H vibrations in the 2940-2840 om-1 region which overlap the hydrogen bonded N-H absorption. However, there do not appear to be any major differences between the solid and solution spectra. (We have been unable to trace crystal structure determinations for either I or III.) In Table 1 the positions of the free N-H and the multiple overlapping sub-bands are given for I, II and III for both solid dispersed in a KBr disc and solution. In general there is agreement to within ~15 cm-l between the present results for II and those of BELLA~ and ROGASCH[4]. Further detailed discussion of the hydrogen bonded absorption is complicated because of the multiplet structure but the band-widths are comparable (I N 650 [la] N. SHEPPARD,Hydrogen Bonding, p. 85. (Editedby D. HADZI) Pergamon,London (1959). [17] H. W. EHRLIOH,Acta Cry&. 18, 946 (1960). [18] D. M. W. ANDERSON,J. L. DUNCAN and F. J. C. ROSSOTTI,J. Chm. Sot. 140(1961).
The infrrt-redspectra and hydrogen bonding of heterocyclicthioamides
1903
Table 1. N-H absorption bands (cm-l) of thiazoline-2-thione(I), benzthiazoline-2-thione(I1) and 4.methylthiwoline-2thione(III) in solid (KBr disc) and solution in carbon tetraohloride II
I Solid
3130 3000 2930 2866 2846 2700
Solution 3420 3130 298O(sh) 2930 2870 2860 -
Solid 3190 3130(sh) 3100 3070 3030 3010(sh) 2960 2880 2840 2760 2730 2060
III Solution 3410 3190(sh) 3110 3070 3036 302O(sh) 2960 2890 2840 2760 2730 2680
Solid
3186 3116(sh) 3086 3036 3000 296O(sh) 2916 2900 2866 2836 2710 266O(sh)
Solution 3410 3186(sh) 3120 3086 3040 3006 2960 2930(sh) 2906 2876 2860(s) 2710 266O(sh)
sh denotes shoulder.
II h) 760 cm-l, III N 800 cm-l). The band intensity of III is about 60% tha.n that of II which is considerably larger than that of I. The approximate band centres axe: I, 3020; II, 2960; III, 2880 cm-l. There have been several attempts to correlake Av(v~ - v~.,.~) with the energy of a hydrogen bond, and it is generally reoognised that such correlations should be used cautiously [19-241. However, for the compounds I, II and III both band intensities and Av indicate that the strength of the hydrogen bonding increases in the order I < II < III. This is in agreement with the dimerization constants (see later) and the conoluaion of ELLISand GRIFFITHS [3] from a study of the effect of hydrogen bonding on the U.V. spectra of these compounds. The hydrogen bonded band centre for III does not difFersignificantly from the band centre for thiopyridone [4] indicating that the N-H . . . S hydrogen bond in III is aa strong as that in thiopyridone, one of the strongest so far recorded [25,26]. cm-l,
greater
Determhation of association constants As 8 consequence of the low solubility of the thioamides I, II and III the choice of solvent for the study of association constants was restricted. Carbon tetrachloride was found to be the most suitable. The ~88 of carbon tetrachloride for such studies has been the subject of much discussion because of the possibility of hydrogen R. M. BADGER and 8. H. BAUJSR,J. Chem. Phyu. 6, 839 (1937). R. M. BADGER,J. Chm. Phys. 8,288 (1940). N. D. COUGESIXALZ, J. Chern.Phyu. 18,978 (1960). E. R. LIFPINCOTJIand R. SCHILOEDEFG, J. Clrem.Phyu. 28,lOQQ (1966). and mcdexdarStructure,p. 413. (Edited by MANSEL H. E. HALUX, Irafra-redSpectrosco2ry DAVIES) Elsevier, London (1963). Chapter 8, Methuen, London (1968). [24] L. J. BELLABSY,Advanca in Infra-red .9pect~o.9copy, [26] B. R. -OLD, Acta CY’y8t. 6, 707 (1963). [26] N. KULEV~KYand W. REINEKE, J. Phys. Chem. 73,3339 (1968). [lQ] [20] [21] [22] [23]
1904
P. J. F.
GRIFFITHS, G.
D. MORGAN and B. Ems
bonding to the chlorine atoms [27-311.In the present study we conclude that hydrogen bonding beween the monomeric form of the thioamides and carbon tetrachloride does not occur because water, a stronger hydrogen-bonding base, does not affect the i.r. spectra [see EXPERIMENTAL].The degree of solvation of the monomer and dimer would not be expected to be the same of course. From the similarity of the i.r. spectra of II in both the solid state and carbon tetrachloride solution it is concluded that the dimeric form is predominant in solution. If only the monomer-cyclic dimer equilibrium 2MBT
= (MBT),
(1)
is involved the stability constant &,[32] is defined by B 20
=
IWBThl [MBT12
and the total concentration, B, of thioamide is given by B = b + 2/9,,b2
(3)
where b is the concentration of monomer [32]. The concentration of monomer was determined by measurement of the apparent molar absorptivity of the free N-H band at ~3410 cm-l. If Beer’s law is obeyed b = E . B/E~ where .a0is the molar absorptivity of the free N-H band and E is the apparent molar absorptivity measured at a stoiohiometric concentration B. Direct determination of a0is not possible because there is still significant hydrogen bonding even at the lowest concentration at which the free N-H absorbance could be determined accurately. Substitution for b in equation (3) and rearrangement of terms yields
so that pa0 may be calculated if a plot of l/e against EB is linear. The l/e against EB plots for I, II and III are linear (Fig. 1). The slopes and intercepts were calculated by the method of least squares and the values of a0 and pzo calculated from these are given in Table 2. This treatment has been used previously by HARRISand HOBBS[33] and VINO(XRADOV [34]. In general, if cyclic dimers and linear oligomers of variable size are formed then a complicated relationship is obtained B = b + 2p2,b2 + IZ npnObn
(6)
[27] V. LIDDEL andE.D.B~~mm,Spectrochim.Acta lo,70(1967). [28] P. HWSKENS, R. HENRY and G. GILLWOT, Bull. Sot. Chtkn. France 720 (1962). [29] D. A. IBBITSON and L. F. MOORE, J. Chem. Sot. (B) 76 (1967). [30] A. BALL md J. L. WOOD, Spectmchim. Acta 24& 1108 (1968). [31] A. N. ~ETOHER, J. .Wqp. Chem. 73,2217(1969). [32] F. J. C. ROSSOTTI and H. ROSSOTTI, The determination of stability cm&an& McGraw-Hill, New York (1961). [33] J. T. HARRIS JR. and M. E. HOBBS, J. Am. Chms. Sot. 76, 1419 (1964).
[34] S. N. VINOURADOV, Can. J. Chews. 41, 2719 (1963).
The i&cl-red spectra and hydrogen bonding of heterooyolicthioamides
41 30
I
50
I
I
70
90
1
I
I
I
I
110
130
150
170
190
tEl(xl03)
Fig. 1. (a) Plot of l/e vs. EB accordingto equation (4).
12 -
II -
IO -
6-
?-
65
/ I 30
34
I 36
I 42
I
I
I
I
I
I
46
50
54
56
62
66
(I /a
x
10-3
Fig. 1. (b) Plot of data for I according to the rearrangedequation (4): 1
60
EB
esB
-=---
%O 60
because of the small slope for these date in (a). (Slope = e. = 192)
1905
1906
P. J. F. GaIFPITHS,G. D. MORUANand B. ELLIS Table 2. Molar absorptivities (Q) and self association constants (j&) Compound
80
I II III
192 f 10 196 & 10 200 f 10
88 f 9 670 f 60 5640 f 600
The values of so and Be0in this table were obtained by a least squares numerical fit to equation (4) and not by a graphical method.
Equation (6) is reduced to an expression similar to equation (4) if cyclic and linear dimers only sre in equilibrium with monomer [35]. Thus, when a linear plot of l/e sg&st EB ia obtained it is not established unequivocally that cyclic dimer only is present. However, from the similkty of the i.r. absorption in the hydrogen bonded region for II both in the solid and solution it appears to be definitely established that cyclic dimers are present in solution. This is an alteration of our previous point of view [3], based on the crystal structure proposed by ZVONKOVA et al. [6], that linear oligomers were present in solution. While the evidence is not so dekite it appears that I and III exist predominantly as cyclic dimers also in solution. The small amounts of open-chain dimer which might possibly be present would also absorb at the monomer frequency. It is unlikely, however, that the correction to ,&, would be greater than about 10% [36]. From the geometry of the molecules and the crystal structure determination [7] for II it is seen that the cyclic dimers have linear N-H . . . S hydrogen bonds and we conclude thet essentially, in carbon tetrachloride solution (26’C, 0.5-l x 10-z M) the equilibrium is between cyclic dimer and monomer. Increase in temperature results in sn inoresse in intensity of the 3410 cm-i bsnd due to free N-H rend&reduction in that of the broad hydrogen bonded N-H band, for each of the three compounds. This indicates that the enthalpy changes on dimerization are negative, if it is assumed that the temperature coefficient of e,, is small [37]. In each of the three compounds, although the hydrogen bonding involves the S II -NH-Cgroup of atoms, the effect of the remsinder of the moleoule is quite pronounced. Steric effects csnnot be involved snd the hydrogen-donor capacity is governed largely by the extent of conjugstion of the heterocyclic ring with the S -NH-(!!group. For thiazolidine-2-thione there is some conjugation with the S atom in the heterocycle lectding to a. slight enhancement of hydrogen bonding [36] R. C. LORDand T. J. Poaao, 2. EZektrodwm. 64,672 (1960). [36] N. KULEVSKY and P. M. F~OEEIJCH, J. Am. Chem. Sot. S&4839 (1967). 1371 G. C. Pnawr~~ and A. L. MCCLZELLILN, TlaeHyhogen Bond, p. 76. W. H. Freeman, San Fran&co and London (1960).
The i&k-red spectra and hydrogen bonding of hetsrocyclic thioamides
1907
capacity. For 4-methylthiazoline-2-thione, however, there is much more extensive conjugation with the double-bond between C4 and C5, with consequent considerable increase in hydrogen bonding capacity. The feebly electron donating methyl group would not be expected to have a significant effect. With benzthizoline-2-thione the S II
extent of conjugation of the double bond between C4 and CS with the -NH-Cgroup is reduced by the conjugation in the six membered ring and there is a concomitant reduction in hydrogen bonding’capacity. This demonstrates the considerable importance of the covalent contribution to the hydrogen bond in these compounds. The N or S methyl derivatives of these compounds can be prepared easily, indicating that chemically they exhibit tautomeric behaviour although normslly they exist predominantly in the thione form [Z, 31. In view of the dual nature of the chemical reactivity of these compounds it is probable that the difference between the energies of the thione and thiol forms is not large so that co-operative proton transfer [38] in the excited state, between the forms V and VI
‘N
I ?
A‘S
i
Q
would be expected to occur to a larger extent than customarily with NH . . . S bonds. In accordance with this view is the fact that the hydrogen bonded band widths are similar, although the positions of the band centre and pa0 values are consistent with the increasing strength of hydrogen bond formed in going from I to III, indicating that the lttrge band-width is not merely a consequence of the usual relationship between AY and vt [39]. Acknowledgements-We are grateful to Messrs. E. WANNELLand J. CANNARDfor experimenti &k3t&IC0. [38] C. G. CANNON,S~ectrochim.Acta 10,341 (1968). [39] G. C. PIXENTELand A. L. MCCLELLAN,2% Hydrogen Bond, Francisco ad London (1960).
p. 94.W. H. Freeman, San