KMROF IODINE-ALKYL SULFIDE COMPLEXES
4017
Nuclear Magnetic Resonance of Iodine-Alkyl Sulfide Complexes'
by E. Thomas Strom, Wilson L. Orr, Brinkley S. Snowden, Jr., and Donald E. Woessner Field Research Laboratory, Mobil Oil Corporation, Dallas, Texas 76921 (Received J u n e 9 , 1967)
Equilibrium constants (Kt's) were determined by high-resolution nmr for iodine complexes of 11 cyclic and aliphatic sulfides in CC14 a t 25". I n the majority of cases the precision is better than 10% and compares favorably with absorption spectroscopy. The Kf's agree reasonably well for the five obtained by conventional absorption spectroscopy. The chemical shifts of the protons in the complexes were tabulated. Substitution of methyl groups for cy protons increases the Kr, but this effect is subject to steric hindrance. The large Kt of the thiacyclopentane-iodine complex is discussed.
It has long been known that iodine and alkyl sulfides form molecular complexes.2 Hastings found that iodine and aliphatic sulfide complexes have large absorptivities in the near-ultraviolet r e g i ~ n . ~Effective absorptivities are a function of the absorptivity of the complex and the position of equilibrium 1 for each sulfide. Equilibrium constants and absorptivities have RSR'
+ Iz E RSR'.Iz
(1)
been evaluated for a number of alkyl sulfides by spectrophotometric Our goals are to determine the effect of complexing on the chemical shifts of the protons, to compare equilibrium constants meaured by nmr with those determined by absorption pectroscopy, and to elucidate the structural features ontrolling the strength of the complex formation.
Results If a proton undergoes a fast exchange between two different chemical species, a single proton resonance is observed a t a position which depends on a weighted average of the position of proton resonance and the relative abundance of the two species.'** For the specific example of the protons CY to the sulfur atom in an alkyl sulfide, we denote the chemical shifts of the a protons in the uncomplexed sulfide and the iodinecomplexed sulfide as 6"s and 6"~, respectively. For a dynamic system in which reaction 1 occurs, the experimentally measured chemical shift of the a protons, 6",,, is given by b",, = 6"cA
+ 6"s(l
- A)
(2)
where A is the mole fraction of iodine-complexed sul-
fide and (1 - A) is the mole fraction of uncomplexed alkyl sulfide. The mole fraction of complexed alkyl sulfide is thus given by A =
6",, 6°C
- 6"s - 6"s
(3)
Remembering that the sum of concentrations of complexed and uncomplexed sulfide must equal the original alkyl sulfide concentration, [RSR' lo, expression 4 for the molar concentration of alkyl sulfide-iodine complex can be obtained. The equilibrium constant
[RSR'*Iz]= A[RSR'Io
(4)
for complex formation, Kr, is thus given by eq 5.
Kf
=
A [RSR'JO
([RSR']' - A[RSR']')([Iz]O - A[RSR']')
(5)
Many previous nmr studies of charge-transfer complexes have dealt with complexes in which the K i s are not large, and it was impossible to determine chem(1) Presented at the 153rd National Meeting of the American Chemical Society, Miami Beach, Fla., April 1967. (2) G. Patein, Bull. Sac. Chim. Paris, 50, 201 (1888). (3) S. H. Hastings, Anal. Chem., 25, 420 (1953). (4) (a) N. W. Tideswell and J. D. McCullough, J . Am. Chem. Sac., 79, 1031 (1957); (b) J. D. McCullough and D. Mulvey, ibid., 81, 1291 (1959). (5) H . Tsubomura and R. P. Lang, ibid., 83, 2085 (1961). (6) M. Tamres and S. Searles, Jr., J. Phys. Chem., 66, 1099 (1962). (7) (a) H. S. Gutowsky and A. Saika, J . Chem. Phys., 21, 1688 (1953); (b) H. 9. Gutowsky, D. W. McCall, and C. P. Slichter, ibid., 21, 279 (1953). (8) H. M. McConnell, ibid., 28, 430 (1958).
Volume 7 1 , A'tcmber 1 9 November 1967
4018
E. T. STROM, W. L. ORR,B. S. SNOWDEN, JR.,AND D. E. WOESSNER
ical shifts for the protons in the complex d i r e ~ t l y . ~ J ~ The shifts for pure complex and the K ( s must then be determined by modified Benesi-Hildebrand plots. 9-11 The Kr's for the sulfide-iodine complexes are large; 0 ,'Ot 0 B 8 hence, the chemical shifts for the protons in the complex are readily determined, and the Kr's are conveniently evaluated numerically. Equilibrium constants in CCl, were determined in the following manner. First, the chemical shifts for the uncomplexed sulfide were measured. These were concentration independent for the ranges studied. I3O Then chemical shifts for the protons in the complex were measured by diluting successively smaller amounts of 0.1 M stock sulfide solution to 1 ml total volume with 0.05 M 1 2 solution. Addition of iodine caused a downfield shift of the proton resonances. When [I;] further addition of iodine no longer caused changes in proton frequencies, the measured proton chemical Figure 1. Plot of K fus. [I4O for three shifts were assumed to be the shifts for the pure comiodine-alkyl sulfide complexes. plex. The basic splitting patterns appeared unchanged by complexing. If more concentrated It solutions were used to determine the shifts for pure complex, ties of the cyclic sulfide-iodine complexes are also reany or all of the following happened: the proton corded in Table I. frequencies never reached a limiting value, peaks Discussion broadened markedly, and new peaks appeared. Reasons for this behavior may include formation of comOne expects from simple considerations of local plexes not obeying eq 1, oxidation to sulfoxides,'2 or diamagnetic currents that complexing of the sulfur other side reactions. Even sulfide solutions conatom should deshield the a protons, and this is what taining excess 0.05 M It showed signs of decomposition is observed experimentally. If we compare methylene protons, the results for di-n-propyl sulfide and thiaafter several hours. The shifts for pure complex, therefore, were always determined within 10 min of cyclopentane indicate that deshielding of p protons is preparation. approximately half that of the a protons. Methyl protons a to the sulfur are deshielded by 0.29Measurements (usually 8-12) were then made for several ratios of [IZj0to [RSR'IO in the range of total 0.31 ppm. This range is relatively narrow when sulfide concentrations from 0.02 to 0.0025 M . Solucompared with the 0.10-0.15-ppm deshielding range tions containing excess sulfide or only small excesses of of the @-methyl protons. In the aliphatic series the Iz gave constant values of proton frequencies for amount of deshielding seems to increase with increasing methyl substitution. This may simply reperiods of a t least 24 hr. Nevertheless, in general, measurements were made within 8 hr. Figure 1 flect the closer approach of the p-methyl group to the shows plots of Kr vs. [I2]O for typical runs. The iodine. A remarkable chemical shift is found for the dideviations appear to be random. isopropyl sulfide-iodine complex. The a-methine proThe Kr's were determined, whenever possible, from ton is shielded on complexation. This should be coma shifts. When a choice was available between two pared to the complex with methylisopropyl sulfide different a proton resonances, that of lower multiplicity was chosen. The Kr for the diisopropyl sulwhere the methine proton is deshielded by -0.1 ppm. fide-iodine complex, however, was determined from the It seems probable that, in the iodine-diisopropyl sulp shifts. Spectra of the aliphatic sulfides were amenable to firsborder analysis, and the determined Kr's (9) (a) M.W.Hanna and A. L. Ashbaugh, J. Phys. Chem., 68, 811 (1964); (b) H.A. Sandoval and M. W. Hanna, ibid.,70, 1203 (1966). and measured chemical shifts are given in Table I. (10) (a) R. Foster and C. A. Fyfe, Trans. Faraday Soc., 61, 1626 Rather broad lines were found for the a protons in (1965); (b) R.Foster and C.A. Fyfe, ibid., 62, 1400 (1966). thiacyclohexane while the spectrum of thiacyclobu(11) H.A. Benesi and J. H. Hildebrand, J . Am. Chem. Soc., 71, 2703 tane was of higher order. This caused a marked de(1949). crease in the accuracy of measurement. The proper(12) T.Higuohi and K.-H. Gensch, ibid., 88, 6486 (1966).
I
t
The Journal of Physical Chemistry
NMROF IODINE-ALKYL SULFIDE COMPLEXES
4019
Table I : Properties" of Iodine-Alkyl Sulfide Complexes in CCL a t 25' Sulfide
Mess MeEtS Me-i-PrS EtzS n-PrzS" i-PrsS t-BuzS Thiacyclobutane Thiacyclopentane 2,2,5,&Tetramethylthiacyclopentane Thiacyclohexane'
dc'
6%
dag
2.054 2,037,; 2,441' 2. 024' 2. 728d 2.474 2.412 2.890
...
...
1.245 1.241 1.234 1.578 1.224 1.380 2.962' 1.914 140,0,*190.7'
2.522
1.807'
...
3.180' 2.755
0.294 0 . 298,' 0,209" 0.314,'00.10d 0.213 0.227 -0.015
...
- dB8 ...
0.161 0.118,' 0.063c
71.1 f 4 . 8 ( 9 ) 136.2 f 15.8(11) 155.4 f 11.8(10) 170.8 f 23.3(8) 168.8 f 17.1 (11) 184.7 i 14.7(12) 159.3 f 8.0(11) 95.8 f 10.6(6) 215.9 f 16.9(11) 78.4 i 5.7(11)
.*.
136.0 f 29.4(6)
0.097 0.118 0.118 0.123 0.139 0.148
...
0.218 0.305
... 0.210
K f (no. of measurements)
'
" Shifts are given in ppm. Errors are standard deviations. Value for methyl protons. Value for methylene protons. Value for the methine proton. ' 6 Y s = 0.998, 6Yc - 678 = 0.006. This is the position of the highest point of a complex envelope rather than a true chemical shift. The chemical shifts for thiacyclobutane are tias = 3.43 and bB8 = 3.17: E. Lippert and H. Prigge, Be?. Bunsenges. Physik. Chem., 67,415 (1963). The y-proton resonance is centered about 1.643.
'
Table 11: Measured Kf's for Iodine-Alkyl Sulfide Complexes a t 25' Sulfide
-%Clr
Mess EtzS Thiacyclobutaned Thiacyclopentane Thiacyclohexane
71.1 f 4 . 8 170.8f23.3 95.8 f 10.6 215.9 f 1 6 . 9 136.0f29.8
KYiacc1,
Fian.heptsnea
71 f 2'
...
...
187.1 f 3.4 90.3 f 1.8 210.7 f 4.8 153.3 f 1.1
87 i 4,' 79.2" 186 2' 148'
*
Kuvn.heptt.nea
. . ~ 168,' 180.4 f 7 . 3 111.4 f 0 . 9 251.4 f 1.9 155.3 f 1 . 5
'
Except where noted from ref 6. Error limits are for the 50% confidence level. From ref 4a. Error is standard deviation. ' Calculated from data in ref 5. d,The value determined in ref 6 in CClr by ultraviolet spectroscopy is 79.0 f 2.7. e From ref 4b. Errors Work of J. D. McCullough cited in ref 6. The value given in ref 4b is 110. are standard deviations?
'
fide complex, the methine protons have a preferred conformation with respect to the iodine-sulfur bond. The Kf's for some of the complexes covered in this study have previously been determined by visible and ultraviolet spectroscopy. Table I1 compares the results from this nmr study with those obtained by absorption spectroscopy. The agreement is reasonable, even excellent for some compounds. Each method has definite advantages in specific cases. Solvent considerations will at times restrict either method. It is experimentally easier to determine a Kt from a line position, as in the nmr method, than from line intensities, as in absorption spectroscopy. Balancing this advantage is the inherently greater sensitivity of absorption spectroscopy. Analysis of the ultraviolet spectra of iodine-alkyl sulfide complexes in certain cases is complicated by the overlapping of the chargetransfer bands with the sulfide absorption. Also, one has to assume that the absorption of the complex follows Beer's law. These difficulties are obviated by
the nmr method. For sulfides where there are large numbers of uncoupled protons, the nmr method seems excellent. However, the accuracy of the nmr method decreases with a decrease in the number of protons contributing to a given resonance and with an increase in the multiplicity of the resonance. When Kf's can be determined by either method, the nmr method may be preferred because of its experimental convenience although the sensitivity is less than that of the absorption spectroscopy method. For donor-acceptor complexes, such as the iodinealkyl sulfide complexes, simple theory predicts that substitution of electron-donating groups about the donor atom should increase the strength of the complex. This prediction is amply realized. Substitution of a single methyl group for one of the hydrogens of dimethyl sulfide results in a 92% increase in the Kf. iMore modest effects ensue from further substitution of methyl groups. I t appears that either the inductive effect from increased methyl substitution is saturated Volume 7 1 . Number 12
November 1967
E. T. STROM,W. L. ORR,B. S. SNOWDEN, JR.,AND D. E. WOESSNER
4020
or the inductive effect is cancelled by increased steric hindrance to complex formation. The results for the cyclic sulfides (vide infra) seem to indicate the latter. The lower Kr for the iodine-di-t-butyl sulfide complex than for the diethyl sulfide complex clearly indicates that steric hindrance occurs in the former compound. Comparison of the thiacyclopentane-iodine complex and the diethyl sulfide complex shows a 26% increase in Kr for the cyclic compound. The results for the five-membered ring may be rationalized in the following manner. When methyl groups are substituted for CY hydrogens, there will be repulsions between the methyl groups and the electron pairs on the iodine molecule in certain conformations of the complex which will tend to destabilize the complex. The situation has been interpreted in terms of front strain or “F train."'^ When the p-alkyl groups are “tied back” in a five-membered ring, these repulsions are minimized and the Kt of the complex is increased. Prototypes for this concept are the trimethylboronamine c ~ m p l e x e s . ~ * J ~ There is a drastic decrease in Kr when every CY position in thiacyclopentane is methylated. If conditions are optimum for complex stability in the thiacyclopentane-iodine complex, the decrease of stability in the 2,2,5,5-tetramethylthiacyclopentane-i0dine complex is understandable. The Kr of the latter complex is about half that of the corresponding di-tbutyl sulfide complex. This difference may reflect the increase in rotational freedom of the methyl groups in the aliphatic sulfide, thus allowing a conformation with stronger sulfur-iodine bonding. More subtle structural factors influencing the strength of the complex in cyclic sulfides will be examined in the future. Furthermore, iodine complexation can be used as an aid in interpretation of the nmr spectra of bicyclic sulfides. The nmr spectrum of I shows on cursory examination the presence of only
I
two protons CY to the sulfur. Iodine complexation, however, indicates the existence of two other CY protons whose resonances fall under the methylene envelope in the uncomplexed compound.16 The Journal of Phusieal Chemistry
Experimental Section The following comlhercially available sulfides, which were found to be free of impurities by gas-liquid partition chromatography (glpc) , were used : diethyl sullide, methyl isopropyl sulfide, diisopropyl sulfide, and di-t-butyl sulfide (Aldrich) ; dimethyl sdfide and thiacyclohexane (J. T. Baker, “Baker grade”) ; di-npropyl sulfide (J. T. Baker, Technical grade); and thiacyclopentane (Eastman). Thiacyclobutane was an API standard sample containing less than 0.05 mole % impurities. The compound was vacuum distilled immediately prior to use, and the distillate showed no impurities by glpc. A glpc analysis showed that methyl ethyl sulfide (3. T. Baker, Technical grade) and 2,2,5,5-tetramethylthiacyclopentane (gift from Bureau of Mines, Bartlesville, Okla.) contained significant impurities. The 2,2,5,5-tetramethylthiacyclopentane was purified by preparative scale glpc while methyl ethyl sulfide was purified through mercuric chloride complexation, according to the procedure of McAllan, et al.” Iodine was Baker Analyzed reagent grade. The carbon tetrachloride was redistilled. Tetramethylsilane (ThIS) was then added to the solvent to make up a 3oj, solution of TATS in carbon tetrachloride. The iodine and sulfides were weighed into separate volumetric flasks and diluted with solvent to volume. The components were mixed directly in the nmr sample tubes which were then covered with Teflon caps. Accurate transfers were made with microliter syringes. In preliminary experiments anomalous broadenings were noted on occasion when acid-cleaned tubes were used; therefore, new sample tubes were used which were subsequently cleaned only with organic solvents. Spectra were measured on a Varian HR-100 nmr spectrometer. An internal lock of TAIS was used. All frequencies were measured on a Hewlett-Packard 522-B electronic counter which has a precision of A 0 . l c. The accuracy of the line position measurements ranged from 1 0 . 1 cps for the highest sulfide concentrations to k0.3 cps for the dilute samples. To make certain that iodine had no effect on the chemical shift of TMS, the chemical shift of the phtons in a 3% cyclohexane solution was determined for both pure CCL and 0.05 M Izin CC4, with TMS as an internal (13) H. C. Browh, H. Bartholomay, Jr., and M.D . Taylor, J. A m . Chem. SOC.,66, 435 (1944). (14) H. C. Brown and M. Gerstein, ibid., 72, 2926 (1950). (15) H. C. Brown and S. Sujishi, ibid., 70, 2878 (1948). (16) W. L. Orr, et al., unpublished results. (17) D. T. McAllan, T. V. Cullum, R. A. Dean, and F. A. Fidler, J . Am. Chem. SOC.,73, 3629 (1951).
DIELECTRIC RELAXATION DUETO CHEMICAL RATEPROCESSES
lock. The measured chemical shifts were identical. The temperature of the probe for the Kr determinations was 25.0 1.0" as measured from ethylene glycol calibrations. Equilibrium constants were calculated from eq 5 on the CDC 1604B computer of the Field Research Laboratory.
*
402 1
Acknowledgments. The authors wish to thank L. E. Nelson for operation of the HA-100 nmr spectrometer and C. H. Calvert for purifying the sulfide compounds. We are grateful to the staff of API Project 48 (Bartlesvile, Okla.) for the sample of 2,2,5,5-tetramethylthiacyclopentane and to Mobil Oil Corp. for permission to publish this work.
On Dielectric Relaxation Due to Chemical Rate Processes
by Gerhard Schwarz M a s P l a m k - I n a t i t u t far physikaliache Chemie, G6ttingen, Germany,' and Institut f a r molekulare Biologie, Biochemie und Biophysik, StbckheirnlBraunachweig, Germany Accepted and Transmitted by The Faraday Society
(March 88, 1967)
The effect of chemical relaxation on the dielectric behavior of dipolar reaction systems is examined. When only small electric field densities are applied, the equilibrium of the over-all react,ion is not perturbed so that the static dielectric constant remains unchanged. Nevertheless, the chemical process may affect the dielectric relaxation by providing an alternate means of orienting dipoles. In principle, this should be measurable in cases where chemical relaxation proceeds at least with about the same rate as the rotation of dipoles. If the reaction rate is much faster than rotational diffusion, a distinct dielectric relaxation effect occurs which reflects directly the chemical relaxation process. Owing to the unfavorably large rotational diffusion coefficients of small molecules in solution, pertinent reaction systems can be expected, first of all, among those systems which involve macromolecular particles. Potential applications of the phenomenon aiming at the determination of chemical rate data are discussed, including even cases which originally could not be studied because of dipole rotation being too fast.
I. Introduction Dielectric polarization of an isotropic and homogeneous medium is described by the expression
P
=
M/V
=
EO(€
- l)E
(1)
relating the electric dipole moment per unit volume, P, to the electric field density, E, which induces it (M = over-all dipole moment, T' = volume, co = 8.854 X 10-14 f/cm). The quantity E represents the (relative) dielectric constant of the system. Owing to the finite rate of formation, the actual value of the dipole moment M will lag behind its equilibrium value, if the field
is changed fast enough. This is of particular significance for periodic fields of suffciently high frequencies. I n such cases a phase shift between M and E will occur, resulting in an energy absorption (dielectric loss). Also the amplitude of M will be changed. Using complex notation, this dielectric dispersion is adequately described by writing E = EOeiut(Eo = amplitude of the field, i = dw = angular frequency, t = time) and introducing a complex dielectric constant e =
E* = E'
- ie"
(2)
(1) Send inquiries t o the author a t this address.
Volume 71, Number 12 November 1967
