NOTES
Feb., 1963 the experimental rate constant which probably increases as the [Brz]/[Ar] ratio increases. This leads to an overestimation of kBr2. Conversely, the use of initial slopes will result in an underestimation of the experimental rate constant which becomes worse with increasing values of [Brz]/ [Ar]. The value of k~~~determined in this work compares favorably with the values determined by Christie, et u Z . , ~ and Burns and Hornig.' It is noteworthy that the value given by Christie, et aZ.,6 was determined from a study of the photochemical reaction between Brz and Clz by a different experimental technique than that used in this work. The agreement with the k~~~values of other workers supports the result for Thus, it appears that the bromine molecule is a much less efficient third body for the bromine atom recombination than is iodine for the iodine atom recombination. Acknowledgment.-We gratefully acknowledge support by the National Science Foundation through Grant No. NSF-G4181. FORMATION COR-STANTS OF MOLECULAR COMPLEXES OF IODIR-E MOXOBROMIDE AND SOJlE AROMATIC HYDROCARBONS BY R. D. WHITAKER~ AND H. H. SISLER Department of Chemzsti y, Unzaersztg of Florzda, Gaznesvzlle, Florida
Received August 10, 1968
It is well known that the thermodynamic stability of 1:1 charge-transfer complexes of the halogens and interhalogens with various electron donors can be correlated with the relative donor-acceptor strengths of the interacting species.2 Complex stability is measured in terms of the magnitude of the equilibrium constant for the system D
+ IX
DIX
Where X indicates an halogen atom such as I, Br, or C1, and D any suitable electron donor. With a given donor, iodine monochloride has been found to form complexes which are more stable than those formed by iodine. This is true regardless of whether the electron donors are relatively strong, such as the heterocyclic amine^,^ or relatively weak, such as the aromatic hydrocarbons. On the basis of the stabilities 05 the trihaljde ions, Scott5 assigned an intermediate position as an electron acceptor to iodine monobromide with respect to iodine monochloride and iodine. Indeed, with pyridine, 2picoline, and 2,6-lutidine3 as electron donors, the stabilities of the resulting complexes indicate that the ability of the halogen molecules to accept electrons decreases in the order IC1 > IBr > I*. However, all of these complexes are relatively stable, with formation constants of the order of from 102 to lo6. A value for the formation constant of the mesitylene-iodine monobromide complex of 1.5 at 25' in carbon tetrachloride was obtained by Blake and Keefer.6 Since this value R'. 8. F. Summer Research Paiticipant. ( 2 ) L. J. Andrews and R. M. Keefer, Aduan. Inory. Chem. Radzochem., 3, 91 (1961). Many pertinent references to the literature are given in thls review. (3) A. I. Popov and R. H. Rygg, J. Am. Chem. Soc., 79, 4622 (1967). (4) L. J. Andtews and R M. Keefer, ibzd., 74, 4500 (1952). ( 5 ) R. L. Scott, zbtd., 75, 1550 (1953). (1)
523
lies between the formation constants for the corresponding 1:1 mesitylene complexes with iodine monochloride and iodine, it would appear that the expected relationship among these three electron acceptors with relatively weak electron donors is preserved. However, it also is true that the mesitylene complexes of iodine monochloride and iodine possess stabilities which are intermediate, respectively, with regard to the higher and lower homologs of benzene. Thus, with an increasing number of methyl substituents on the donor molecule, the stabilities of the resulting charge-transfer complexes with both iodine monochloride and iodine increase. This observation has been explained in terms of the inductive effect of additional substituent group^.^ The present work was undertaken with the purposes: (1) to determine whether iodine monobromide occupies its expected intermediate position as an electron acceptor, with regard to iodine monochloride and iodine, toward an homologous series of aromatic hydrocarbons and (2) to determine whether the charge-transfer bands, the maximum absorptivities, and the stabilities of the complexes of iodine monobromide with this homologous series vary in the same manner as do these corresponding properties for complexes of iodine monochloride and of iodine. Experimental Materials .-Iodine monobromide was prepared and purified by the method of Popov and Skelly.' The spectrophotometric properties were in exact agreement with literature valuesS of 492 mp for the absorption maximum and of 400 cm.-l mo1e-I for the maximum absorptivity. Spectro grade carbon tetrachloride was used as the solvent without further purification. The aromatic hydrocarbons were from Eastman Kodak Co., White Label grade, with the exception of p-xylene, which was Phillips 66 Reaearch grade. Some of the hydrocarbons which are liquids and were obtained from previously opened bottles were routinely dried and distilled before me. However, subsequent experiments using material from freshly opened bottles without further purification gave results which were consistent with the previous determinations. The solid hydrocarbons were used without further purification. Spectrophotometric Determinations.-Spectra in the region 600-270 mp were determined using a Cary Model 14 recording spectrophotometer. Absorption cells 1 cm. in thickness were used. Stock solutions of iodine monobromide and the particular hydrocarbon under study were prepared just prior to use. The solutions were mixed immediately before the spectrum was determined. The hydrocarbon concentrations in the resulting solutions were in the range 0.1-1 M , and the iodine monobromide conto 5 X lob4 44. A centrations were in the range 1 X hydrocarbon solution of the same concentration as the hydrocarbon in the solution under study was used as the reference. It was found that the iodine monobromide maximum a t 492 mp shifted toward smaller wave lengths in every case. Likewise, a strong charge-transfer band in the near ultraviolet, where the iodine monobromide absorption is quite small, was noted in every case and was ascribed to the formation of a 1: 1 complex. Although all spectra were determined within 5 or 6 min. from the time the solutions were mixed, some of the systems were subject to undesirable side reactions. Hexaethylbenzene was particularly troublesome in this respect, there actually occurring an apparenf red-shift of the iodine monobromide maximum to the region 520 mp, presumably resulting from the formation of free iodine. After the general position of the charge-transfer band had been located for those systems with which this difficulty was encountered, the immediate range of maximum absorbance was scanned using freshly prepared solutions. In this manner, the spectra could be determined within about 1 min. from the time of mixing the iodine monobromide and hydrocarbon solutions and this difficulty was largely eliminated. All determinations were made a t 24". (6) J. H. Blake and R. M. Keefer, ibid., 1 7 , 3707 (1955). (7) A. I. Popov and N E. Skelly, tbzd., 7'7, 3723 (1955).
524
NOTES
Treatment of Data.-The equation used by Benesi and Hildebrand8 to treat spectrophqtometric data has been shown9 t o be adequate under the proper conditions. These conditions include the requirements that the concentration of the donor is in large excess, that only the complex absorbs a t the wave length under consideration, and that only a 1: 1 complex is in equilibrium with free donor and acceptor molecules. These conditions were met by the experimental conditions employed in this study. We have chosen to use a slight modification of the original equation of Benesi and Hildebrand, as suggested by Rose and drag^,^ and t o use their general method of treatment.10pl1 We used the relation
where K-l is the reciprocal of the formation constant for the equilibrium, IBr D @ DIBr (D represents any hydrocarbon donor), [IBrIoand [D]oare the initial molar concentrations of the respective species, A is the experimentally determined absorbance a t a given wave length, and E is the absorptivity of the complex a t that wave length. The intersections of the straight-line plots of K-' v8. assumed E values gave values for K which were averaged. Four different sets of measurements involving different concentrations of iodine monobromide and hydrocarbon were carried out for each hydrocarbon-IQr system studied. The value of K was determined using three different wave lengths-the wave length corresponding t o the maximum absorbance, and wave lengths 4-8 mp below and above the maximum. All the values for K thus determined were averaged and the average deviation computed. Two methods were employed to obtain the maximum absorptivity of the complexes. (I) The method of Rose and Drag0~81~ when applied to data a t the wave length corresponding to the absorption maximum gave directly values for maximum absorptivity. These values were averaged and the average deviation computed. (11) Using the over-all average value for K , based on the three wave lengths, the concentrations of the complex a t the different initial concentrations of iodine monobromide and hydrocarbon were computed. Using the relation A = E ID .IBr], the maximum absorptivity was calculated. The probable error in E was computed on the basis that the average deviation of K represents the probable error in K , The results obtained by both methods are reported since there is no way to decide which method is superior. The comparison of the values obtained by the different methods serves as a further check on the precision of the data.
+
Results and Discussion Table I summarizes the results which were obtained in this study. Except for the systems involving benzene and toluene, the lue for emax is about the same regardless of which method of calculation is employed. Even for these two systems, the two values for Emax fall within the limits of the precision of the data. It is readily obvious that the larger deviations occur for the less stable complexes. For purposes of comparison with the corresponding iodine monochloride and iodine systems, it is perhaps better to use the values of emex derived by method I, since this treatment corresponds more nearly to the Benesi-Hildebrand treatment. Table I1 presents the pertinent data for iodine and iodine monochloride systems. It is clear from a comparison of formation constants that for a given donor, the expected order of stability is realized. For each of the systems studied, the iodine monobromide complex is more stable than the corresponding iodine complex but less stable than the corresponding iodine monochloride complex. The wave lengths of the charge-transfer bands for the iodine monobromide complexes also fall between the respective (8) H. A. Benesi and J. H. Hildebrand, J . A m . Chem. Soc., 71,2703 (1949). (9) N. J. Rose and R. S. Drago, ibid., 81, 6138 (1959). (10) R. S. Drego and N. J. Rose, i b i d . , 81, 6143 (1959). (11) R. S. Drago, R. L. Carlson, N. J. Rose, end D. A. U'enr, ibid., 85, 3572 11961).
VOl. 67 TABLE I FORXATIOS CONSTANTS FOR IBr-AROMATIC HYDROCARBON COMPLEXES (CCL, 24') Xmax
Donor
K , (l./mole)
Benzene 0.44 f 0.07 Toluene .51 f .07 o-Xylene .78 f .09 m-Xylene .78 f .05 p-Xylene .67 f .06 Mesitylene 1.47 f .03" Hexamethylbenzene 4.3 f .5 Hexaethylbenzene 0.34 i .07 Naphthalene .63 f .04 Reference 6 gives the value
(mp)
286 298 308 310 297 321 356
7
e
-
x
(method I)
X 10-8(method 11)
1 2 . 1 f 2 . 6 9 . 6 3Z 0 . 2 11.0 f 1 . 0 9 . 6 f .9 8 . 9 f 0 . 6 9 . 0 3 ~.6 10.3 f . 4 1 0 . 5 f . 5 9 . 1 f .3 9 . 0 1 . 5 10.5 3Z .1 10.4f .1 6 . 1 f .4
6.23Z . 3
362 1 1 . 3 & 1 . 8 1 1 . 0 2 ~ . 7 353 5.0 f 0 . 3 5 . 2 f . 2 1.5 l./m.
TABLE 11" FORMATIOR. CONSTANTS FOR HALOGEN-AROYATIC HYDROCARBON COMPLEXES (CCl,, 25") -----I%
-IC1
systems-
systems----
KO
KO
Xmnx
€ma=
Xmax
fmax
(l./mole)
(mp)
X 10-8
(l./mole)
(mp)
X 10-8
Benzene 0.15 292 Toluene .16 302 o-Xylene .27 316 m-Xylene .31 318 p-Xylene .31 304 Mesitylene .82 332 Hexamethylbenzene 1.35 375 Hexaethylbenzene 0.13 378 Saphthalene 0.25 360 a Data taken from ref. 4.
16.4 16.7 12.5 12.5 10.1 8.8
0.54 0.87 1.24 1.39 1.51 4.59
282 288 298 298 292 307
8.1 8.0 7.9 9.2 6.5 7.9
334
4.0
340 341
6.6 3.9
8.2 16.7 7.6
22.7 1.24 1.39
wave lengths of the charge-transfer bands for the iodine and iodine monochloride complexes. The maximum values for the absorptivities of the complexes likewise is seen generally to follow this same patterntlie one exception being the mesitylene complex. A Consideration of the change of the stability of the iodine monobromide complexes with increasing methyl substitution on the benzene ring shows generally the same trends as do the iodine and iodine monochloride complexes. Thus all three acceptors form more stable complexes with hexamethylbenzene than with the other hydrocarbons studied. However, as has been pointed outj2the stronger electron acceptor apparently shows a greater relative increase in its formation constant with increasing donor strength than will a weaker electron acceptor. The K , values for the hexamethylbenzene and the benzene complexes of iodine are in the ratio of 9:1, whereas the corresponding ratio of K , values for the iodine monochloride complexes is 42: 1. We might expect that the corresponditlg ratio of K , values for iodine monobromide would lie between these ratios. However, the ratio is actually about 10: 1. This smaller increase in stability than might be expected may arise from steric factors. However, an alternate explanation is simply that the K O value for iodine monochloridehexamethylbenzene may be somewhat in error. The possibility that this reported value is too large is strengthened by the fact that the ratios of the KOvalues for toluene, the xylenes, and mesitylene m7ith the respective K , values for the benzene complexes are not far different for any of tlie three halogen molecules.
KOTES
Feb., 1963
It is interesting to note that there is an apparent inversion in the stabilities af the complexes of the xylene isomers with iodine monobromide as compared with iodine monochloride. The order of electron donor strength of these isomers toward iodine monobromide conforms more closely with that found by McCaulay and Lien12 with HF-BFa-methylbenzene systems than does the apparent order of electron donor strength as measured by the stabilities of the corresponding iodine monochloride complexes. However, the differences in stability are so small and the precision of the data is such that any attempt to generalize on this point see:ms to be unwarranted. (12) D. A. SIcCaulay and A. P. Lien, J . A n . Chem. Sor., 7 3 , 2013 (1961).
ROTATION ABOUT THE CARBON-NITROGEN BOXD I V SOME AROMATIC AMINESa BY I. YAMAGUCHI'~ AND S. BROWNSTEIN
HO-r
Hc
H6
NHz
Fig. 1.-Proton
resonance spectrum of o-nitroaniline.
Replacing the hgdrogens by methyl groups might be expected to increase the barrier to rotation in orthosubstituted amines because of enhanced steric effects. However, down to 173'K. no change was detected in the shape of the peaks due to the N-methyl protons in ortho-methyl, broimo, or nitro substituted N,N-dimethylanilines. Chemical shifts for these and other substituted amines to low field of tetramethylsilane are listed in Table I. Because of steric interaction with
Divzszon of Applied Chemistry, National Research Council, Ottaua, Canada Received August 19, 1982
Resonance interaction between the amino group and the benzene ring in aromatic amines is a maximum when the molecule is coplanar. For this reason the barrier to rotation about the benzene-nitrogen bond might be higher than for an aliphatic carbon-nitrogen bond. An especially favorable case for observing this carrier would be o-nitroaniline. Both substituents contribute to the resonance interaction with the aromatic ring and there is evidence of hydrogen bonding between the nitro group and the amino The proton resonance spectra of this compound and of onitro-N,N-dideuterioaniline are shown in Fig. 1. It is apparent that the resonance absorption of the hydrogens bonded to nitrogen lies under that for protons A, B, and C. This peak also is unusually broad. Since the spectra were obtained in a dilute acetone solution a t room temperature it is unlikely that broadening is due to intermolecular hydrogen exchange. If the two protons are non-equivalent due to hindered rotation about the C-N bond, a four line AB spectrum uTould be expected. The individual lines could be broadened, to give the envelope which was observed, either by motional averaging due to rotation about the carbonnitrogen bond a t a rate comparable to the separation between the individual lines in the absence of rotation or by yuadrupolar broadening of the N14-H spin coupling. Unfortunately, solubility considerations severely limit low temperature work on this compound so it is not possible to distinguish between these possibilities. The chemical shifts, in p.p.m. to low field of acetone, and spin couplings are listed below. They were obtained byasecond-order perturbation treatment. Since these chemical shifts are rather solvent dependent, 7 values seem inappropriate. 6~ = 4.90 6~ = 5.23 6c =
4.48
JAB = 8.5 JBC = 8.0 J c x = 8.6
JAG = 1.6 JBX = 1.3 JAX =0
6x = 5.91 (1) Issued as N.R.C. No. 7194; (b) National Research Council Postdoctorate Fellow, 1962-1963. (2) H. B. Kleirens and J. R. Platt, J . A m . Chem. Soc., 71, 1714 (1949). (3) A. E. Lutskii and V. T. Alekseeva, J . Gen. Chem. USSR, 29, 2957 (1959). (4) J. C. Dearden and W. F. Forbes, Can. J . C i ~ e n .38, , 1837 (1900).
HY
TABLE I PEAKPOSITIONS OF THE N,N-DIZIETHYL GROUP' Substituent
Neat
None m-NH2
2 53 2.56
m-OH
..
o-Br p-Br O-CHa m-CHs p-CHa p-NO
2.59
0-NO2
m-N O2
..
2 47 2.61 2.58
..
2.75
..
..
p-KO2 p-CHO .. a Determined a t 25'. Saturated solution.
---CSz C0non.b
4.3 3.7 3O 3.2 5.4 4.5 4.1 5.5
lo 5.6 6.1 lC 7 in CDCl,
eoln.--Position
2.86 2.82 2 78 2.73 2.85 2.60 2.84 2.80 3.14 2.80 2.98 3.12 3 05
Concentration in mole per cent.
the ortho substituent the N-dimethyl group is twisted out of the plane of the benzen ring. If this twisting angle is large it is reasonable to assume a pyramidal configuration about nitrogen rather than planar.2 Averaging of the magnetic environment of the K-methyl groups will occur rf there is inversion about the nitrogen as well as by rotation about the plane of the aromatic ring. Therefore, a low barrier to magnetic averaging of the N-methyl groups in these compounds is easily rationalized. On the basis of studies involving C14-labeled methyl groups it has recently been reported that one of the Nmethyl groups in o-methyl-N-trimethylaniliniumiodide is not equivalent t o the other tw0.5 This non-equivalence had a half-life of a few hours, but was not observed in the para-methyl or unsubstituted derivative. However in the proton resonance spectrum only a single peak has been observed by us for the N-trimethyl group of the o-methyl and oo substituted compounds. The chemical shifts to HzO for solutions of these salts in DzO are listed in Table 11. It is unlikely that chemically non-equivalent methyl groups attached to nitrogen in either of these compounds would have an identical magnetic environment. There is a difference in chemical shift for the K-trimethyl group in the o-bromo and o-methyl compounds and a differ(6) L. Otvos, F. Dutka, and H. Tudos, Chem. Ind. (London). 818 (1962).
