COMPLEXES OF IODINE WITH TETRAMETHYLUREA AND TETRAMETHYLTHIOUREA energy prohibits the penetration of xenon into the channels. Since the xenon atom is spherical, there is no possibility of passage with a preferred orientation. However, the above theoretical calculation, as admitted by Kington and Laing, assumes the silicate rings to be rigid and takes no account of the amplitude of vibration, concerning which there appears to be no information available at yesent. Moreover, in this system an increase of 0.1 A in the radius of the ring can decrease the value of the activation energy by 67 kcal/mol and the value of 2.55 used in the calculation is the minimum distance from the hexagonal axis. Nevertheless, the explanation offered for the higher value of the surface area of beryl using xenon at -78” compared with that using argon at -196” is to be considered at present hypothetical. I t is tentatively suggested that an explanation similar to the one advanced in the case of muscovite and talc is valid for beryl also.
Appendix I n the calcuolations of activation energy, values of 3.92 X lO-Z4 A3 and 20.92 X 10-30 cgs unit for the
2129
polarizabilities and diamagnetic susceptibilities, respecda and tively, of the oxygen ion12 and 4.01 X 71.5 X 10-30 cgs unit for the corresponding quantities of the xenon a t 0 m ~ ~ swere ~ 4 used. The internuclear separation re was assumed, as advocated by MacLeod and Kington, 15 to be given by the sum of half the kinetic diameter of xenon, r0/2”’, where ro is the ecpilibrium separation of two xenon atoms,16viz., 4.46 A, and the (Goldionic radius of oxygen is taken to be 1.32 Schmidt). The values of r were obtained from the square root of the sum of the squares of the radius of the ring and of the distance along a line perpendicular to the plane of the ring and passing through its center. (12) L. Pauling, Proc. Roy. Soc., A114, 181 (1927). (13) J. A. Beattie and W. H. Stockrnayer, “Treatise on Physical Chemistry,” Vol. 11, D. Van Nostrand Co., Inc., Princeton, N. J., 1951,p 297. (14) P. W.Selwood, “Magnetochemistry,” Interscience Publishers, Inc., London, 1956. (15)A. C. MacLeod and G . L. Kington, Trans. Faraday SOC.,55, 1799 (1959). (16) E.A. Mason and W. E. Rice, J. Chem. Phys., 22, 843 (1954).
Molecular Complexes of Iodine with Tetramethylurea and Tetramethylthioureal by Robert P. Lang Department of Chemistry, Quincy College, Quincy, Illinois
62601 (Received November $0, 196’7)
Absorption spectrophotometric studies in the near-ultraviolet and visible spectral regions have been made on the iodine complexes of tetramethylurea and tetramethylthiourea, in both n-heptane and dichloromethane. KO,AH”, and AS” have been determined for both complexes in both solvents. A comparison is made of the K Ovalues ,determined for both the uv and the visible spectral regions for the tetramethylurea-iodine complex. Thermodynamic evidence for the location of the donor site, in both the tetramethylurea and the tetramethylthiourea complexes with iodine, is discussed. The relative effect of a polar and a nonpolar solvent on the complex-formation equilibria and thermodynamic data is considered and rationalized in terms of Mulliken’s charge-transfer (CT) theory. The spectral characteristics of the C T band have also been determined for both complexes. The effect of solvent on the CT-band position is considered and also rationalized in terms of Mulliken’s C T theory. Data on the “blue-shifted” visible iodine band are also reported.
Introduction Although extensive spectrophotometric studies have been made on molecular complexes of iodine with various amides, in dichloromethane2“ and carbon tetrachloride,Zb complete information is still lacking on the spectral characteristics of the charge-transfer band of an amide-iodine complex. Information on the long-wavelength side of the chargeetransfer band in the ultraviolet spectral region, along with thermodynamic data, has been reported for theiodine complexes with acetamide and N,N-dimethylformamide, in dichloro-
methane. 2s Since the spectral characteristics for both the charge-transfer band and the “blue-shifted” visible iodine band have been determined for a great number of iodine complexes covering a relatively large range of donor strength3 it seemed worthwhile to obtain this infOi-mation for an amkb-iodine complex. (1) Presented before the Division of Physical Chemistry at the 155th National Meeting of the American Chemical Society, San Francisco, Calif., April 1968.
@)(a) H. Tsubomura and R. P. Lang, J . Amer. Chem. SOC.,83, 2085 (1961); (b) R. S. Drago, D. A. Wens, and R. L. Carlson, ibid., 84, 1106 (1962). Volume ‘78, Number 6 June 1968
2130 The thermodynamic and electronic spectral characteristics of several thioamide-iodine complexes in di~hloromethane,~ chloroform,6 and carbon tetrachloridea have been determined. However, no spectrophotometric determinations have been reported either for the thioamide-iodine complexes or for the amideiodine complexes, in a relatively more “inert” solvent such as a saturated hydr~carbon.~Drago has recently made a study of solvent effects on the thermodynamic characteristics of an amide-iodine complex.* He points out the possible hazards of comparing thermodynamic data for complexes studied in polar solvents with that of complexes studied in “inert” solvents. Consequently, it was thought t o be of some interest to determine the thermodynamic and electronic spectral characteristics of an amide-iodine complex and the structurally analogous thioamide-iodine complex in both an inert solvent and a polar solvent. Dichloromethane and n-heptane were chosen as the polar and inert solvents, respectively. Commercial availability, solubility, and absorption spectrum in the visible and near-ultraviolet spectral regions dictated the choice of tetramethylurea and tetramethylthiourea.
Experimental Section Materials. Matheson Coleman and Bell spectroquality n-heptane and dichloromethane were used without further purification. Resublimed iodine and reagent grade tetramethylurea, obtained from Matheson Coleman and Bell, were also used without further purification. Tetramethylthiourea (E( & K Laboratories) was recrystallized once from absolute ethanol. SpectrophotometricDeterminations. Absorption spectra in the visible and near-ultraviolet spectral regions were measured by a Beckman double-beam (DB) ratiorecording spectrophotometer equipped with a Sargent linear-log (SRL) recorder. A pair of matched Beckman 1-cm silica “U” rectangular cells with ground-glass stoppers were employed for all the spectroscopic measurements. The Beckman DB spectrophotometer has a thermostatible cell compartment through which thermostated water was circulated from a WAC0 (Wilkins-Anderson Co.) heating and cooling constanttemperature bath. Measurements were made at 10, 15, 20, 25, and 30” with the sample temperatures maintained constant to better than *0.5”. Temperatures lower than 10” could not be employed because of the condensation of water vapor on the cell windows (a dry-nitrogen purge kit cannot be used with the DB spectrophotometer). Because of the low boiling point of dichloromethane, temperatures above 20” were not employed for this solvent. Temperatures above 30” were not employed for the n-heptane solutions in order to obtain comparable data for the standard enthalpies of complex formation in n-heptane and dichloromethane. For all the equilibrium-constant determinations, a The Journal of Physical Chemistry
ROBERT P. LANG series of solutions with a constant initial iodine concentration and variable initial tetramethylurea and tetramethylthiourea concentrations was employed. The donor and acceptor concentrations were chosen so that all absorbance readings were in the range 0.300.95. Treatment of Data. The equilibrium constants for the tetramethylurea-iodine complex in both n-heptane and dichloromethane were evaluated by using eq 3 of ref 4 (usually called the Scott modificationg of the Bensi-Hildebrand equationlo) with absorbance data from the peak of the charge-transfer band. At this wavelength, with the concentrations employed, there was a small absorption due t o tetramethylurea and a very small absorption due to iodine. Since the concentration of tetramethylurea was considerably greater than that of iodine and, therefore, also of the complex, the absorption due to tetramethylurea could be directly subtracted from the total absorbance, as if all the tetramethylurea were uncomplexed, without introducing any significant error. The very small absorption due to uncomplexed iodine was first neglected in the calculation of K,. Then, employing this value of K,, a correction for uncomplexed iodine was made and a new value for K , was calculated. This correction for iodine absorption turned out to be of negligible importance, as the two values for K , were within experimental error of each other. The blue-shifted visible iodine band of the tetramethylurea-iodine complex, in both n-heptane and dichloromethane, was overlapped considerably by the visible iodine band of uncomplexed iodine. Consequently, the following equation1’ (which takes this uncomplexed iodine absorption explicitly into account) was employed to evaluate K , from absorbance data in the visible spectral region
where [DIois the initial concentration of donor, e, is called the apparent molar extinction coefficient of the (3) See, for example, Table I11 of H. Tsubomura and R. P. Lang, J . Amer. Chem. SOC.,83, 2085 (1961). (4) R. P. Lang, ibid., 84, 1185 (1962). (5) K. R. Bhaskar, R. K. Gosavi, and C. N. R. Rao, Trans. Faraday SOC.,6 2 , 29 (1966). (6) R. J. Niedzielski, R. S. Drago, and R. L. Middaugh, J . Amer. Chem. SOC.,8 6 , 1694 (1964). (7) For example, the relative inertness of n-heptane toward iodine,
compared to dichloromethane and carbon tetrachloride, can be estimated by comparing the location of the absorption maximum of the visible iodine band for these latter two solvents relative to n-heptane.3 (8) R. S. Drago, T. F. Bolles, and R. J. Niedzielski, J . Amer. Chem. SOC.,88, 2717 (1966). (9) R. L. Scott, Rec. Trav. Chim. Pays-Bas, 75, 787 (1956). (10) H. A. Bensi and J. H. Hildebrand, J . Amer. Chem. SOC.,71, 2703 (1949). (11) This equation could be called the Scott modification of the Ketelaar equation: J. A. A. Ketelaar, C. van de Stolpe, A. Goudsmit, and w. Dzcubas, Rec. T T ~ V .Chim. Pays-Bas, 71, 1104 (1952).
COMPLEXES OF IODINE WITH TETRAMETHYLUREA AND TETRAMETHYLTHIOUREA
2131
Table I : The Thermodynamic Characteristics of the Tetramethylurea-Iodine and the Tetramethylthiourea-Iodine Complexes AHo, kosl/mol
KO,
Solvent
+Heptane Dichloromethane Carbon tetrachloridea n-Heptane Dichloromethane Carbon tetrachloride* Chloroformc
l./mol (20°)
Donor, Tetramethylurea 19 i 2 4.6 i 0.5 6.5 i.0 . 2 (25') Donor, Tetramethylthiourea 13,000 It 1100 49,000 i 5100 8,000 i 200 (25') 13,560 (25")
ASO,
eu
-5.0 i0.4 -3.3 i0.6 -4.4 f 0.2
-11 i 1
-9.9 & 1.2 -9.0 i 1 . 6 -10.5 i 0.5 -9.5
-15 i4 -9i5
-83~2
R. J. Niedzielski, R. S. Drago a R. L. Middaugh, R. S. Drago, and R. J. Niedzielski, J. Amer. Chem. SOC., 86, 388 (1964). K. R. Bhaskar, R. K. Gosavi, and C. N. R. Rao, Trans. Faraday SOC.,62, 29 and R. L. Middaugh, ibid., 86, 1964 (1964). (1966).
acceptor, E, is the molar extinction coefficient of the complex, €1, is the molar extinction coefficient of uncomplexed iodine, and K O is the molar concentration equilibrium constant for 1:1 complex formation. ea = A/[IzIo, where A is the experimental absorbance reading and [I2I0is the initial concentration of acceptor. This equation, which as presented here is valid for a path length of 1 cm, reduces to the Scott modification of the Benesi-Hildebrand equation when the molar extinction coefficient for uncomplexed iodine equals zero. A plot of [D]o/(E*- a*)us. [DIoshould yield a straight line of slope l / ( ~- €1,) and intercept 1/Ko(e0 - a,), There was no absorption by tetramethylurea in the visible spectral region, so the reference cell contained only pure solvent. For the solutions used to study the tetramethylthiourea-iodine complex, at the maximum of the charge-transfer band, the initial concentration of tetramethylthiourea and of iodine only differed by about a factor of 10. Consequently, eq 2 of ref 4 was employed in the evaluation of KO. There was only a negligible absorption due t o iodine in this spectral region, at the concentrations employed. However, there was a small absorption due t o tetramethylthiourea which was directly subtracted from the total absorbance, as was done in the case of tetramethylurea. The standard enthalpies and standard entropies of complex formation were evaluated in the usual way from a plot of log K , us. 1/T. However, since the temperature range was rather small, the estimated error in the value of the equilibrium constant and the experimental uncertainty in the temperature were utilized in determining the estimated error for the enthalpies and entropies.
Results and Discussion The thermodynamic characteristics of the tetramethylurea-iodine and the tetramethylthiourea-iodine complexes, in n-heptane and dichloromethane, are listed in Table I. The spectral characteristics of the
charge-transfer band for both of these complexes, in both solvents, are listed in Table 11. The thermodynamic values listed in Table I were determined from ~~
Table I1 : The Spectral Characteristics of the Charge-Transfer Band of the Tetramethylurea-Iodine and the Tetramethylthiourea-Iodine Complexes (mnxs
Solvent
Xmaxl
mr
1. mol-] om-1
D,
Ail,/,,
om-]
n-Heptane Dichloromethane
Donor, Tetramethylurea 9,000 7300 268 282 7,000 8100
%-Heptane Dichloromethane Chloroforma
Donor, Tetramethylthiourea 338 35,000 4500 328 36,000 5300 334 36,600 5300
f
debyes
0.28 0.24
4.0 3.8
0.68
7.0 7.6 7.7
0.82
0.84
a K. R. Bhaskar, R. K. Gosavi, and C. N. R. Rao, Trans. Faraday SOC.,62, 29 (1966).
the peaks of the charge-transfer bands for both complexes, in both solvents. For the tetramethylureaiodine studies, the concentrations of tetramethylurea and iodine were in the range 0.02-0.2 and (1.5-3.5) X M , respectively. In the case of tetramethylthiourea, its concentration was about (1.0-5.0) X 10-4 M and that of iodine was (2.5-3.0) X 10-5 M . Isosbestic points were obtained in the visible spectral region for both complexes, in both solvents, which supports the interpretation that these systems involve a dynamic equilibrium between uncomplexed iodine and one form of complexed iodine. Comparison of K O Data from Both Uv and Visible Spectral Regions. The values of K , (20"), for the tetramethylurea-iodine complex, determined from spectroscopic data in the visible spectral region were 13 k 1.5 (n-heptane) and 1.5 f 0.3 (dichloromethane). These values are a little smaller than the K , values Volume 7.8, Number 6 June 1S68
ROBERTP. LANG
2132 determined from the charge-transfer band in the ultraviolet (Table 1).l2 Since the extinction coefficient of the charge-transfer band is several times larger than that of the blue-shifted visible iodine band, more concentrated solutions have to be employed for the visible spectral studies than for the ultraviolet studies. The iodine concentration was about four times greater and the tetramethylurea concentration was about twofold greater for the visible studies than for the ultraviolet studies. Solutions with greater tetramethylurea concentrations were not employed because their spectra missed the isosbestic point. Trotter and Hanna13have pointed out the difficulties that can arise from a Benesi-Hildebrand-type method of determining equilibrium constants and extinction coefficients for molecular complexes when conditions of solution ideality do not exist. Since the concentrations of uncomplexed iodine, uncomplexed tetramethylurea, and the complex were all greater for the visible spectrophotometric studies than those employed for the ultraviolet studies, data obtained from the latter studies should correspond more closely to conditions of solution ideality than data obtained from the former studies. Therefore, the value of K , based on uv data should be closer to the thermodynamic equilibrium constant than that based on visible data. Thermodynamic Evidence for Location of Donor Site. Infrared studies indicate that the donor site in amideiodine complexes is at the oxygen atom.14 Both nmr6 and infrared5 studies support the conclusion that the sulfur atom acts as the electron donor in thiourea-iodine complexes. Arguments based on thermodynamic data have been used to help establish the donor site in iodine complexes with acetamide and thi~acetamide.~ The thermodynamic data obtained in the present study can also be utilized to help determine the donor site. The structural environment of the nitrogen atoms in tetramethylurea and tetramethylthiourea is essentially that of a tertiary amine, with a decreased nitrogen lone-pair donor strength due to the intramolecular resonance enhancement of the oxygen and sulfur atoms, respectively. Tertiary amine -iodine complexes, in n-heptane, have AH"'s of about -12 kcal/mol.'6 This value is a little greater than the A H o of the tetramethylthiourea-iodine complex and substantially greater than that of the tetramethylurea-iodine complex. If a nitrogen atom acted as the donor site in these iodine complexes, one would expect the AH" t o be a little less than that of a tertiary amine-iodine complex, owing to the intramolecular resonance. While the AH"'s determined in the present study make the nitrogen atom donor site hypothesis plausible for the tetramethylthiourea-iodine complex, they certainly tend to make the hypothesis unsatisfactory for the tetramethylurea-iodine complex. This hypothesis would seem to imply that the intramolecular resonance produced a less than 10% reduction in donor strength for tetramethylT h e Journal of Physical Chemistry
thiourea, whereas it produced a greater than 50% reduction for tetramethylurea. Furthermore, one would not expect such a great discrepancy in resonance effects between tetramethylurea and tetramethylthiourea. I n addition, it is believed that resonance effects are greater for thiocarbonyl derivatives than for the analogous carbonyl derivatives. l6 The acetamide and thioacetamide complexes with iodine, in dichloromethane, have AHO's of -4.6*& and -8.2 kcal/rn01,~ respectively. The nitrogen atom environment in these donors is essentially that of a primary amine, and primary amine-iodine complexes have AHO's of about - 7 kcal/mo1.l5 The AHO's of the acetamide and thioacetamide complexes with iodine are in much closer agreement with those of the tetramethylurea and tetramethylthiourea complexes with iodine than one would expect if a nitrogen atom were the donor site in these complexes, in view of the difference in donor strength between primary and tertiary amines toward iodine. Furthermore, how could one explain the fact that the thioacetamideiodine complex is stronger than a primary amine-iodine complex? A comparison of the AHO's of the tetramethylurea and tetramethylthiourea-iodine complexes, in n-heptane, with those of the diethyl ether and diethyl sulfideiodine complexes, in n-heptane, tends to support the conclusion that the donor sites are at oxygen and sulfur for the former two iodine complexes as well as €or the latter two. The AHO's for the diethyl ether and diethyl sulfide-iodine complexes are - 4.217 and - 8.9 kcall mol,'* respectively. The resonance effect should tend to make the oxygen and sulfur atoms in tetramethylurea and tetramel hyIthiourea better electron-donor sites than the oxygen and sulfur atoms in diethyl ether and diethyl sulfide. The fact that the AH"'s for the tetramethylurea- and tetramethylthiourea-iodine complexes are greater than for the diethyl ether and diethyl sulfide-iodine complexes is consistent with the hypothesis that oxygen and sulfur act as the donor sites for the former two iodine complexes. Also the difference in A H o between the diethyl ether- and diethyl sulfideiodine complexes (4.7 kcal/mol) and that between the (12) In order to determine if these apparent differences in K, were due to a difference in the method of calculation, the visible data were treated by the same method used with the uv data (Scott-modified BH equation). With iterative recycling t o correct for absorption due to uncomplexed iodine, the resulting KOvalues were essentially the same as those obtained from the Scott-modified Ketelaar equation. (13) P. J. Trotter and M.W. Hanna, J . Amer. Chem. Soc., 88, 3724 (1966). (14) C. D. Schmulbach and R. S. Drago, ibid., 82, 4484 (1960). (15) H. Yada, J. Tanaka, and S. Nagakura, BUZZ.Chem. Soc. Jap., 33, 1660 (1960). (16) K. R. Bhaskar, S. K.Bhat, A. S. N. Murthy, and C . N.R. Rao, Trans. Faraday SOC.,62, 788 (1966). (17) (a) AM.Brandon, M. Tamres, and S. Searles, J . Amer. Chem. SOC.,82, 2129 (1960); (b) M. Tamres and M. Brandon, ibid., 82, 2134 (1960). (18) M. Tamres and S. Searles, J . Phys. Chem., 66, 1099 (1962).
COMPLEXES OF IODINE WITH TETRAMETHYLUREA AND TETRAMETHYLTHIOUREA tetramethylurea- and tetramethylthiourea-iodine complexes (4.9 kcal/mol) tends to support the conclusion that oxygen and sulfur lone pairs act as the donor sites for all of these iodine complexes. Thermodynamic and spectral data have recently been reported for iodine complexes with acetone and thiocamphor in cyclohexane.l6 For the acetoneiodine complex, K , (25") and AH" are 0.8 l./mol and -2.5 1 kcal/mol, respectively. These values are somewhat smaller than the corresponding data for the tetramethylurea-iodine complex, which is what one would expect if the carbonyl oxygen atom acts as the donor site for both of these complexes. The thermodynamic data for the thiocamphor-iodine complex are 95 l./mol (25") and -11.0 1 kcal/mol for K , and A H " , respectively, This K , value is substantially smaller than that of the tetramethylthiourea-iodine complex, whereas the AHO's for both of these complexes are essentially the same when one considers the estimated experimental uncertainty. Also, this K , value is consistent with the hypothesis that the thiocarbonyl sulfur atom acts as the donor site for both the tetramethylthiourea- and the thiocamphor-iodine complexes, because one would expect resonance enhancement t o make the K , value of the former complex somewhat greater than that of the latter. Furthermore, this conclusion is supported by the fact that, although there is a substantial difference in K , values for these two complexes, the thiocamphor K Ovalue is comparable to that of the diethyl sulfide-iodine complex, just as the analogous acetone K , value is comparable to that of the diethyl ether-iodine complex. Finally, if one considers the magnitude of the estimated experimental uncertainty in AH" values, the A H o for the thiocamphor-iodine complex could quite possibly be somewhat smaller than that of the tetramethylthiourea-iodine complex; therefore, it may very well be that the correct AHO's parallel the trend observed in the K,'s. Solvent Efect on Thermodynamic Datu. Thermodynamic and spectral data for the iodine complexes of benzene,19 diethyl ether,lg and diethyl sulfide2(' in the gas phase have recently been reported. For the benzene-iodine complex, which is the weakest complex of the three,21 both K , and AH" are smaller for the gas phase than for solution in an inert solvent. The K , values for the gas-phase diethyl ether and diethyl sulfide-iodine complexes are both smaller relative to an inert solvent, whereas the AH" data are comparable for both gas and inert-solvent systems for both of these iodine complexes. Furthermore, the relative difference in K O values between vapor and solution phases decreases as the strength of complex formation increases. According to the results of the present study of the tetramethylurea-iodine complex, both K , and AH" are smaller for the more polar solvent, with Drago's results for carbon tetrachloride being intermediate between the n-heptane and dichloromethane values
*
*
2133
(Table I). For the N,N-dimethylacetamide-iodine complex Dragos has obtained a similar decrease in both K , and AH" in going from carbon tetrachloride to dichloromethane, with intermediate results for the donor solvent benzene. Also, for the ethyl isothiocyanate-iodine complex, which is comparable in strength to an amide-iodine complex, KlaboeZ2reports a decrease in K , in changing solvent from n-heptane t o dichloromethane. I n contrast to the previously discussed results, the present study shows that for the tetramethylthioureaiodine complex, the K , value is greater for the more polar solvent. The A H o values are comparable in both solvents, according to the present study, and are, therefore, of little value as a criterion for determining how complex strength is affected by solvent polarity. Both TamresZoband Personzshave recently pointed out that the product KCeocan be determined more precisely than either term alone. However, in the present case the product Koecis also greater for the more polar solvent (the E, values are essentially the same for both solvents as indicated in Table 11). Kobinata and NagakuraZ4have obtained very large increases in K , for two amine-iodine complexes in going from the inert n-heptane solvent to the donor solvent dioxane. They also report increases in E , along with the K , increases. Other workers5report very little variation in K , values for some dialkylthiourea-iodine complexes between the solvents carbon tetrachloride and chloroform. However, these workers report a marked increase in K , for the tetramethylthioureaiodine complex in chloroform compared to carbon tetrachloride (Table I). KlaboeZ6has recently studied the Raman spectra of the thiourea-iodine complex in both the nonpolar solvent benzene and the polar donor solvent acetonitrile.26 The Raman spectra results indicate a stronger donor-acceptor interaction in the polar solvent. For the strong triphenylarsine-iodine complex, KlaboeZ7also reports a stronger donor-acceptor interaction in the polar solvent dichloromethane than in the nonpolar solvent carbon tetrachloride. These results can be rationalized within the context of Mulliken's charge-transfer theory of molecular (19) F. T. Lang and R. L. Strong, J . Amer. Chem. Soc., 87, 2345 (1965). (20) (a) J. M.Goodenow and M. Tamres, J. Chem. Phgs., 43, 3393 (1965); (b) M.Tamres and J. M . Goodenow, J . Phys. Chem., 71, 1982 (1967). (21) R. M. Keefer and L. J. Andrews, J. Amer. Chem. SOC.,77, 2164 (1955). (22) E. Plahte, J. Grundnes, and P. Klaboe, Acta Chem. Scand., 19, 1897 (1965). (23) W. B. Person, J . Amer. Chem. SOC.,87, 167 (1965). (24) 8. Kobinata and S. Nagakura,.ibid., 88, 3906 (1966). (25) P. Klaboe, ihid., 89, 3667 (1967). (26) (a) A. I. Popov and W. A. Deskin, ihid., 80, 2976 (1958); (b) W. B. Person, W. C. Golton, and A. I. Popov, ihid., 8 5 , 891 (1963). (27) E. Augdahl, J. Grundnes, and P. Klaboe, Inorg. Chem., 4, 1475 (1965).
Volume 7.8, Number 6 June 1968
2134
ROBERT P. LANG
complex formation.28 According t o Mulliken’s valence-bond formulation, the ground (N) and excited (V) states, responsible for the experimentally observed intense long wavelength charge-transfer (CT) band, of a molecular complex can be approximately represented by the wave functions ~ N ( D * Az) a\k(D, A)
+ b\k(D+-A-)
(2)
\kv(D+*A-) m b * 9 (D,A)
- a*\k(D+-A-)
(3)
and where a = a* and b = b*. P(D, A) and \k(D+-A-) represent the “no-bond” and “dative-bond” functions, respectively. The excited state ‘kv(D+-A-) is often referred to as the charge-transfer state. The effect of solvents on complex formation equilibria depends upon whether the solvent has a greater interaction with the complex or with the donor and acceptor. The solvent effect reversal in K , values, obtained in the present study, indicates that the tetramethylthiourea-iodine complex is stabilized to a greater extent in a polar solvent than the tetramethylurea-iodine complex. For weak complexes the wave function coefficient a (in eq 2 ) is much greater than the wave function coefficient b, with the a / b ratio decreasing as the donor-acceptor interaction increases. Consequently, the b/a ratio should be much greater for the tetramethylthiourea-iodine complex than for the tetramethylurea-iodine complex. Therefore, the former complex should have a greater dative chargetransfer contribution to the ground state than the latter complex, resulting in the observed greater stabilization in a polar solvent. The tetramethylureaiodine complex must have a relatively low polarity, since a polar solvent does not favor complex formation and the polarity of the tetramethylthiourea-iodine complex must be rather high, since dichloromethane stabilizes the complex more than the uncomplexed donor and acceptor. The Charge-Transfer (CT) Band. The reversal in CT band maxima, observed in the present study, can also be rationalized in terms of Mulliken’s theory. According to the theory of solvent effects,2gfor a polar solute dissolved in increasingly polar solvents, a “red shift” is predicted when the ground state is less polar than the excited state, whereas a “blue shift” is expected when the ground state is more polar than the excited state. From eq 2 and 3, one can see that when the coefficient a is much greater than the coefficient b, the ground state (N) consists mainly of the no-bond function and the excited state (V) consists mainly of the dative-bond function, resulting in a relatively high polarity for the excited state and a low polarity for the ground state. When the coefficient b is greater than the coefficient a, the relative polarity of the ground and excited states is reversed, producing a ground state that is more polar than the excited state. The Journal of Physical Chemistry
For the intermediate strength tetramethylureaiodine complex, one would expect the coefficient a to be substantially greater than the coefficient b, resulting in a lower polarity for the ground state, of the chargetransfer band, than for the excited state. This could account for the observed red shift of the chargetransfer band of this complex with increasing polarity of the solvent. Since the tetramethylthiourea-iodine complex is a rather strong complex, one would expect the coefficients a and b t o be comparable. The high polarity of the ground state of this complex is also supported by the previously discussed effect of a polar solvent on the equilibria, and one might even expect the coefficient b to be a little greater than the coefficient a. This would result in a ground state, for the charge-transfer transition, that is more polar than the excited state and could account for the observed blue shift of the charge-transfer band with increasing polarity of the solvent. Analogous solvent effects on the CT band have been obtained for the ethyl isothiocyanate22 and triphenylarsine-iodine ~omplexes.~’In going from a nonpolar to a polar solvent, the CT band of the former complex is red shifted, whereas the CT band of the latter, much stronger, complex is blue shifted. From Table I1 one can see that there is a considerable increase in CT band intensity for the tetramethylthiourea-iodine complex compared with the tetramethylurea-iodine complex. There is a comparable increase in CT band intensity for the diethyl sulfideiodine complex relative t o the diethyl ether-iodine complex (both in Both pairs of iodine complexes have greater values for the extinction coefficient, oscillator strength, and transition dipole moment for the sulfur-donor complexes than for the oxygen-donor complexes. I n addition, both pairs of iodine complexes have smaller half-widths for the stronger sulfur-donor complexes than for the weaker oxygen-donor complexes. The Blue-Shifted Visible Iodine Band. The blueshifted visible iodine band of the tetramethylureaiodine complex is overlapped on the long-wavelength side by the visible iodine band of uncomplexed iodine. This overlap was accounted for by resolving the experimental total absorbance into “free” and “complexed” iodine components (through use of the previously determined K,). The resulting blue-shifted visible iodine band of the tetramethylurea-iodine complex, in n-heptane, had a wavelength maximum at about 460 mp and an extinction coefficient at this wavelength of about 1200 1. mol-’ cm-’. For the tetramethylthiourea-iodine complex, in (28) (a) R. S. Mulliken, J . Aner. Chem. SOC.,74, 811 (1952); (b) R. S. Mulliken, J. Phys. Chem., 56, 801 (1952). (29) H.H.Jaffe and M. Orchin, “Theory and Applications of Ultraviolet Spectroscopy,” John Wiley and Sons, Inc., New York, N. Y., 1962,pp 186-195.
2135
cis AND trans ISOMERS OF A MONOSUBSTITUTED FORMANILIDE n-heptane, this long-wavelength overlap by uncomplexed iodine is considerably less, owing to the greater blue shift of the visible iodine band of this complex upon complex formation. However, the blue-shifted visible iodine band of this complex is overlapped on the short-wavelength side to some extent by the CT band. An absorption maximum appears at about 440 mp, with an extinction coefficient of about 3600 1. mol-' cm-'. From the shape of the absorption spectrum, a reasonable estimate of about 25% overlap by the C T band (as an upper limit) could be made. This would make the extinction coefficient of the blue-shifted visible iodine band at the 440-mp absorption maximum about 2700 I. mol-I cm-l. For the tetramethylthiourea-iodine complex, in dichloromethane, the blue-shifted visible iodine band appears only as a long wavelength shoulder on the CT band, without a distinct peak. It is interesting t o compare these data with the corresponding values for the diethyl ether and diethyl sulfide-iodine complexes, in n-heptane. The wavelength maxima and extinction coefficients are 462 mp and 950 1. mol-' cm-I for the former complex and 435 mp and 1960 1. mol-' cm-I for the latter complex. For both pairs of iodine complexes the stronger sulfur donor produces a greater blue shift and a greater intensity for the visible iodine band of the complex relative to the oxygen donor. The extent of the blue
shift in wavelength is of about the same magnitude for both pairs of complexes. The extinction coefficients of the urea-thiourea pair are a little greater than for the ether-sulfide pair. I n a recent paper, h9cCull0ugh~~ reports thermodynamic and electronic spektral data for the strong selenacyclopentane-iodine complex, in carbon tetrachloride. This complex is similar in donor strength to the tetramethylthiourea-iodine complex, with a KO (25') of 2200 l./mol and a A H o of -11.1 kcal/mol. It is interesting to note that the spectral characteristics of both the CT band and the blue-shifted visible iodine band are also similar to those of the tetramethylthiourea-iodine complex. For the CT band, the wavelength maximum is at about 322 mp and the extinction coefficient is equal t o 39,600 1. mol-' cm". The blue-shifted visible iodine band peak is at about 428 mp, with an extinction coefficient of 3350 1. mol-' cm-l. It would be very interesting to have the corresponding data for the tetramethylselenourea-iodine complex.
Acknowledgments. This research was supported, in part, by the National Science Foundation under Grant GP-5470. The author wishes to thank Mr. David Aalbers for his assistance with the laboratory work. (30) J. D. McCullough and A. Rrunner, Inorg. Chem., 6, 1251 (1967).
Proton Magnetic Resonance and Infrared Studies of the cis and trans Isomers of a Monosubstituted Formanilidel by T. H. Siddall, 111, W. E. Stewart, and A. L. Marston Savannah River Laboratory, E . I . d u Pont de Nemours and Go., Aiken, South Carolina 93801
(Received November 81,1367)
Pure cis and trans isomers of o-methylformanilide were isolated. When the amide dissolved in chloroform, an equilibrium cis: trans isomer ratio of 3: 1 was established. The kinetics of the cis-trans equilibrium a t several temperatures in deuteriochloroform were studied by proton magnetic resonance. Infrared studies of the cis and trans isomers were made, and band assignments for the amide vibrations are given.
Introduction Amides have been the subject of many investigations by means Of proton magnetic resonance* Many Of these investigations have been devoted to a study of the rate of rotation around the carbonyl-to-nitrogen bond (amide bond)*2 Recently t b isomers Of unsymmetrically substituted amides have been separated.*
When a separated isomer is brought to the appropriate temperature in solution, its Pmr signals decay as the (1) The information contained in this article was developed during the course of work under Contract AT(07-2)-1 with the U.8. Atomic Energy (2) J. W. Emsley, J. Feeney, and L. H. Sutcliffe, "High Resolution Nuclear Magnetic Resonance," Pergamon Press Ltd., London, 1966. (3) A. Mannschreck, Tetrahedron Lett., 1341 (1965). Volume 79,Number 6
J u n e 1068
