KOTES
2486
the emission bands certainly implies that the fluorescence spectrum (Fig. 4) belongs to acridine (Fig. 1). Moreover, the above argument also implies that the acridine mas free from hydrogen bonded water, since, if that were not the case, the addition of ethanol would not be expected to result in further red shifts in the absorption and emission spectra. Finally, we wish to note the mirror image relationship which exists between the fluorescence (Fig. 4) and the portion of the absorption spectrum (Fig. 1) corresponding to the longest wave length transition. This relationship is more easily seen by referring to the absorption spectrum for acridine hydrochloride (Fig. 3) in which more of this transition is exposed. The mirror image relationship would not be expected to exist if emission occurred from a lower lying n,n* state. Thus, the possibility of fluorescence from an n,n* state to the red of the 'La state can be eliminated. The only basis for expecting acridine not to exhibit a ffuorescent emission would be that conversion of the excited molecule to a triplet state or internal conversion were so highly probable that the normal fluorescence could not compete. Under the circumstance, one might expect t o see a reasonable intensity of phosphorescence. However, such is not the case, since experiment has shown the acridine phosphorescence to be very weak. 2b Nevertheless, the argument could be kept intact if a reason for high conversion to a triplet state could be provided, accompanied by a strong internal conversion mode out of the triplet. Such a reason would exist if acridine exhibited a lowest singlet excited state of n, n* character. This usually results in complete conversion to the triplet state. However, recent experimental work by Coppens, Gillet, Nasielski, and DoncktlO places the n,x* state a t 360 mp, %ell above the 0-0 vibrational level of the 1L, state of acridine. Also, this investigation points to the fact that no n,n* state is below a n,n*one. In addition, Bennett1' performed lifetime measurements on the emission of acridine in a water solution and (16) G. Coppens, C. Gillet, J. Nasielaki, and E. V. Donckt, Spectrochim. Acta, 18, 1441 (1962). (17) R. G . Bennett, Rev. Sci. Instr., 81, 1275 (1960).
Vol. 67
observed two exponential decays of 3.8 X see. and 4.3 X sec. He attributed, via chemical evidence, the longer lifetime to the acridine cation and the shorter one to the acridine molecule. Thus, both the negative arguments and the positive results of Bennett give strong indication that an intrinsic fluorescence should exist for acridine. This is, of course, in agreement with the results of this investigation. Conclusion The acridine molecule has a moderately strong inherent fluorescence. We feel that substantial evidence has been provided to justify its identification as such. It has been shown that the intrinsic fluorescence of acridine can be associated with the longest wave length (T+ n*) electronic transition. This and more specifically the mirror image relationship allows us to conclude that acridine does not exhibit a lowest lying n,n* electronic excited state. The extreme sensitivity of acridine to solvent impurities, mentioned above, has been clearly illustrated in this study by the observed formation of complexes a t low temperature between acridine and a solvent impurity other than water. Both the shifted absorption spectra and fluorescence spectra of these complexes have been presented. Such data in the case of acridine hydrochloride strongly suggest that these complexes are not formed via hydrogen bonding type interaction. The acridine hydrochloride case further suggests that for the solutions of acridine in EPA at 77'K., the observed complex specie involves the acridine-ethanol hydrogen-bonded system. This would mean that the three-component complex (acridine, ethanol, and solvent impurity) results from two different weak interactions. Parallel type complexes have been reported earlier. The complexes appear to involve peroxides. Acknowledgments.-We wish to thank Dr. R. W. Harrell of E. I. Dupont for first suggesting that a further study of acridine would be interesting and for helpful discussions concerning it.
NOTES CONFORMATIONS OF ACETANILIDE AND N-METHYLACETANILIDE1 BY H. BRADPORD THOMPSON AXD KARBN M. HALLBERG
relative position of the carbonyl and the aromatic ring should be indicated readily. Results of this study and of EL similar investigation for methylacetanilide are reported in Table I.
Alfred Nobel Science Laboratories, Gustavus Adolphus College, 9t. Peter, Minnesota
TABLE I ELECTRIC DIPOLEMOMENTS
Received April 16, 1968
The question of the stable conformation(s) in acetanilides has interested several investigators.2.B There is evidence that one planar or nearly planar structure predominates.* By comparison of the electric dipole moment of acetanilide with p-bromoacetanilide, the (1) The aid of the National Science Foundarion through a basio research grant and through a n undergraduate research participation program supporting K. M. H. is gratefully acknowledged. We wish to thank Dr. Coluroba Curran, who, as referee for this paper, made a pumber of helpful suggestions. (2) J. W. Smith, J. Chen. Soc., 4700 (1961). (3) R. A. Russell and H. W. Thompson, Spectrochim. Acta, 8,138 (1956).
Compound
lI.p.,OC.
Aoetanilide p-Bromoaoetanilide h'-Methylacetanilide p-Bromo-N-methylaoetadide Bromobenzene a
Slope for (E
See ref. 7 and 8.
- $)/(e
114' 167.5-167.8 104.0-104 4 97.5- 97.8
+
2 ) (712 See ref. 2.
Y E l e o t r i o momentsSo This study Smithb 102.5 3.88 3.65 136.2 4.47 4.36 87.2 3.57 36.8 2.32 15.0 1.48
+ 2) us. molar concentration.
The acetanilide case has been studied using electric dipole moments' by Smith,2 whose data are included
NOTES
Nov., 1963 for comparison in Table I. Smith chose to analyze the results in terms of group moments for the carbonyl (from acetone) and amine (from aniline) groups. From these and the observed moments for acetanilide and related compounds he deduces the angle of rotation about the S-GO bond. Several assumptions are necessary jn this approach: the geometry about the nitrogen must be the same in amine and amide; the moment associated with the lone pair on the amine nitrogen must be retained in the amide; and the resonance structure I1 must make a very limited contribution. Alternatively, we have chosen to assume that the moment
H
CH3
\ / C--iS / \
0
CsH6
I
H
CH3
\
+/
C=iY
/
\ CC"
-0'
I1
of p-bromoacetanilide differs from that of acetanilide by the vector addition of 1.49 D. in the direction of the C-Br bond. We have assumed the acetanilide moment to be a t an angle of about 130' from the N-CO bond, and nearly parallel to the C=O bond, as shown in Fig. la. This angle may be deduced from the bond moments employed by Kotera, Shibata, and Sone.4 Since the C-0 moment is certainly the largest coiitributor in acetanilide, this orientation is probably not greatly in error. The moment according to Smith would be a t a significantly smaller angle and would not lie in the OCN plane. However, even this moment would not qualitatively alter the argument to follow. If acetanilide has 1,he planar structure commonly assumed for the amide group,5 or if (like formamide6) it is only slightly nonplanar, then the moment of p-bromoacetanilide should be predicted by one of the two vector drawings in Fig. l a ; the structure with the ring cis to the oxygen best matches the data. This is the conformation found by Brown and Corbridge to be stable in the solid.' Since the solid struclxre involved intermolecular hydrogen bonding, it mas not obvious that the same conformation would be most stable in solution. Brown and Corbridge found a dihedral angle of 37' between the plane of the ring and that of the acetyl; this appears to arise from steric repulsion and may well be retained in solution. However, unlike the rotation about the OC--N bond assumed by Smith, the rotation in the crystal is about the CeH6-K bond; that is, a tilting of the ring. This would have little effect on t)he electric moment. The differerice between Smith's moments and those reported here may be due to the choice of solvent. Smith used benzene; we chose dioxane, since monosubstituted amides may form hydrogen-bonded dimers in benzene. Hydrogen bonding to the solvent seemed preferable, since this should vary little between the compounds being compared. Smith's data would lead to the same conclusion as ours when similarly analyzed. The reverse situation clearly holds in N-methyl(4) A. Kotera, S. Shibata, and K. Sone, J . A m . Chem. S O C . , ? ~ 6183 , (1955); the exact direction iriredicted aould depend on the conformation: 1 2 5 O for the hydrogen trans t o the carbonyl, 137' for the ring trans. ( 5 ) L. Pauling, "The Nature of the Chemical Bond," 3id Ed., Cornel1 University Press, Ithaca, N. Y . , 1960, p. 281. (6) C. C. Costain and J. M. Dourling, J . Chem. Phys., 32, 158 (1960). (7) C. J. Brown and D. E. 0.Corbridge, Acta Cryst., 1, 711 (1964).
24:87
Fig. 1.-Vector addition of component electric moments: a, p-bromoacetanilide; b, p-bromo-N-methylacetanilide. Solid arrows are component moments; broken arrows are predicted molecular moments.
acetanilide. From Fig. lb, our experimental mometnt of 2.32 D. corresponds to the structure with ring and oxygen trans. This result appears surprising a t first; however, infrared studies3 on monosubstituted formamides and acetamides have indicated that the tendency of an N-methyl group to take up the position cis to the oxygen is stronger than that of an E-phenyl. Experimental Electric moments %-ere determined in dioxane solution by Guggenheim's method.* Procedure and solvent purification have been described p r e v i o ~ s l y . ~ Commercial acetanilide and p-bromoacetanilide were recrystallized from benzene. N-Methglacetanilide was prepared by addition of acetyl chloride to chilled methylaniline in benzene. The unchanged acet) 1 chloride was distilled off, the crystalline product dissolved by heating, and the solution filtered and cooled. A sample of the solid product was recrystallized repeatedly from benzene and used for the electric moment measurement. A portion of the remainder was brominated by the procedure of Ashley and Berglo and the product recrystallized from petroleum ether. Each solid v a s dried for several hours in a vacuum oven before solutions were prepared. Melting points of the samples used are included in Table I. Commercial bromobenzene was distilled twice in a small fractionating column and the product dried over CaC12; b.p. 155.3-155.5" (740 mm.), n Z 61.5575, ~ da6,1.4909. (8)E. A. Guggenheim, Trans. Faraday hoc., 45, 714 (1949). (9) H. B. Thompson, I,. Eberson, and J. V. Dahlen, J . Phys. Chem., 66, 1634 (1962). (10) J. N. Ashley and
S. Berg, J . Chem.
S o c . , 3089 (1957).
&4CID-BASEREACTION3 IX FUSED SALTS. THE DICHROMATE-BROMATE REi4CTION' B Y F. R. DUKEd N D JAMESSCHLEGEL Instatute for Atomic Research and Department of Chemastry Iowa State Unacerszty, Ames, Iowa Recezved Apral 29, 1969
Some equilibrium studies of Lewis acids and bases in fused alkali nitrates have been made in which Crz07-2 was used as an acid, and others in which Br03- was used as a base. In the first case, Crz0,-2 was allowed to react with nitrate to form NOz+.2 I n the second case, a heavy metal ion was added with BrO3- and the decomposition of the metal-bromate complex was stutlied.3 When a divalent metal ion reacts with bromate, the metal-bromate complex, MBr03+, is formed arid slowly decomposes. This was postulated on the ba& that the rate of decomposition varies with metal ion used. However, if one uses Cr207-2 as an acid, no complex is formed, and BrOz+ is the probable equilibrium product with bromate present. Thus, a direct COM(1) Contribution No. 1279; work mas performed in the Ames Laboratory of the U.S. Atomic Energy Commission. (2) F. R. Duke and 31.Iverson J Phys Chem., 6 2 , 417 (1958). (3) F. R. Duke and W. W. Lawrence, J . A m . Chem. S o c . , 88. 1271 (1961).
