64
M. A.
Battiste, M. W.
Couch,
and R. Rehberg
Electronic Spectra and SolvatOchromism of the pPolyphenylyltropylium Ions and a Comparative Study of the Cyclopropyltropylium Ion M. A. Battlste, M. W. Couch,+ and R. Rehberg Departments of Chemistry and Rad;ology, University of Florida, Gainesville, FlorMa 326 10 (Received June 2 1, 1976) Publication costs assisted by the Department of Radiology, University of Florida
Three p-polyphenylyltropylium salts (phenyltropylium fluoroborate (l),p-biphenylyltropyliumperchlorate (2), and p-terphenylyltropylium(3)) were prepared and their electronic spectra examined in several solvents. A bathochromic shift of the long wavelength maximum arose in this series upon successive addition of p-phenyl groups to the tropylium cation residue. This shift, which was markedly solvatochromic, was attributed to an aryl substituent-to-tropylium ion ring intramolecular charge-transfer transition. The tropylium ion was shown to be a superior probb of the electron donor properties of the cyclopropyl ring in cyclopropyltropylium hexachloroantimonate (5).
Introduction In connection with earlier attempts to determine the efficacy of the tropylium ion as a sensitive intramolecular probe of electronic donor properties of an attached substituent in monosubstituted tropylium ions we prepared and examined the electronic spectra of the p-polyphenylyltropylium ion salts phenyltropylium tetrafluoroborate ( l ) ,p-biphenylyltropylium perchlorate (2), Ar@
X-
3.Ar=
m, m,
X = C104 X=c1O4
4,Ar = C H 3 O e , X = C l 0 4
5, X = SbCt6 or c104
Ring protons"
Cation (solvent) Tropyliume
H3C SbC16 m H3C 6
and p-terphenylyltropylium perchlorate (3). In this series the successive addition of p-phenyl groups to the tropylium residue produced the expected bathochromic shift of the long-wavelength maximum, this band extending into the visible region of the spectrum in 2 and 3. On the basis of theoretical and the effect of substituents on the position and intensity of the long-wavelength band, this band has been attributed to an aryl substituent-totropylium ring intramolecular charge-transfer transition by us3 and othersa4Both Harmon's recent disclosure of previously unpublished data from Dauben's laboratory4 and our observations on the isoelectronic relationship of the p-polyphenylyltropylium series to the classic ppolyphenyl series prompted us to describe in full our spectral results for the series 1-3 including a study of their solvatochromic behavior. In addition we provide an illustration of the utility of the tropylium cation as a superior probe of the electron donor properties of the cyclopropyl ring in the cyclopropyltropylium ion (5). Experimental Section Synthesis of the p-polyphenylyltropylium ions involves the preparation of the 7-monoaryltropilidene from the corresponding aryllithium reagent and tropylium fluoroborate, thermal isomerization to the 3 isomer, and finally hydride abstraction from the 3-monoaryltropilidene by means of trityl fluoroborate (to give 1) or perchlorate (to give 2,3, and 4).$1° Salts 1,2, and 4 have been previously ~ynthesized.~ p-Terphenylyltropylium, prepared in three steps from p-terphenyl-4-yllithiumll and tropylium The Journal of Physical Chemistry, Vol 8 1, No. 1, 1977
Tropylium
I ~ (. C H* , C N, )
9.2 9.15
2c(CF, COOH)
9.02
I , A r = o , X=BF4 2, A r =
P-@ x-
TABLE I: Proton Chemical Shifts for Some Tropylium Cations
3c3d (CF,COOH) 5, X =
SbC1,- (CD,CN)
6,X = SbC1,- (CD,CN)
-8.6 8.89
9.2
Aryl
Alkyl
protons"
7.86 7.68 7.79 7.43 7.69 7.53 7.31 2.55 (aCH) 1.87 (PCH) 1.55 (PCH) 3.65 (CH) 1.56 (CH,)
a In ppm downfield from TMS. Tetramethylsilane as external standard. Tetramethylammonium fluoroborate as internal standard. Tropylium proton signal partially masked by solvent. e Reference 16.
fluoroborate, was obtained in approximately 80% overall yield, mp 279-282 "C (dec). Anal. Calcd. for CZ5Hl9C1O4: C, 71.68; H, 4.57; C1, 8.46. Found: C, 71.47; H, 4.77; C1, 8.25. Additional proof of structures 1-3 was afforded by the reaction of the monoaryltropylium ions with hydrogen peroxide to give the corresponding monoarylben~enes.~~' Thus 1, 2, and 3 gave biphenyl, p-terphenyl, and pquarterphenyl, respectively. Cyclopropyltropylium hexachloroantimonate (5, X = SbCls-) (mp 123 "C. Anal. Calcd. for C10H~~Sbc16: C, 25.79; H, 2.38. Found: C, 25.68; H, 2.40.) and the less stable perchlorate (5, X = ClO,), mp 82 OC, were prepared by treatment of 7-cyclopropyltropilidene12(synthesized by the reaction of cyclopropyl magnesium bromide13 with 7-tropyl ethyl ether14)with trityl hexachlor~antimonate~~ and trityl perchlorate,6 respectively. The corresponding isopropyltropyliumhexachloroantimonate (5), mp 112 "C, was prepared in a similar manner. Results and Discussion In the NMR spectra of 1, 2, and 3 (Table I), the tropylium protons appeared in the same region as the unsubstituted ion. The 1.90-ppm downfield shift for the unsubstituted tropylium ion with respect to benzene has been attributed to decreased electron density at the in-
65
pPoiyphenylyltropyiium Ions
TABLE 11: Ultraviolet Absorption Maxima (nm) of Monoaryl and Monoalkyl Tropylium Ions in Various Solventsa R@
Tropylium salt
CH,CN
CH,Cl,
1
367(17 600) 269(15 700) 23lC 414(14 200) 275(sh 7950) 247(16 000) 432(17 400) 273(28 100) 432.5(24 500) 274(8700) 232c 327(12 300) 247(29 800) 274(14 400) 225(32 200)
387(18 1 0 0 ) 274(13 100) 232c 449(14 800)d
2 3 4
5 6
7
257(14 800)d -483-49OblC 282b 463(18 700) 279(sh 3780) 238(15 700) 335(16 600) 254(30 900) 277(13 900) 234.5(35 850)
Solute incompletely soluble in this solvent.
a E values in parentheses. cells used.
-,..
361 34
/f /
237b 442b
448(21 1 0 0 )
260b 473b 283b
258629 700) 480 275b
457(21 800) 271(4860)
456(19 600) 271 (7400)
Was not accurately determined.
V ,a,g
Ar
No. C6H5
I 08
I O
12
10-cm
/
0 12
261
CF,COOH 386(15 800) 270(9500)
TABLE 111: Observed Frequencies ( v ), Calculated Transition Energies (Am ) of Monoaryl Tropylium Cations, and Ionization Potentials of the Donor Hydrocarbons (I,,)
/
30
CHCl, 384b
14
16
1.8
Am
Figure 1. Observed frequencies (vmax)of monoaryl tropylium ions and HMO energy differences (Am).
dividual carbon atoms and a subsequent deshielding of the protons in the positively charged seven-membered ring.6 Hence the slight but steady upfield shift of the tropylium protons with each additional p-phenyl group must correspond to some reduction of the positive charge in the seven-membered ring. In other words, delocalization of charge into the phenyl rings in the ground state does occur, if only to a slight extent. Similar arguments hold for cation 5 although it is clear that cyclopropyl electron release is superior to phenyl in the ground state for this cationic system. The downfield shift for the CY and 0protons of the cyclopropyl ring in 5 is consistent with the modest electron demands of the tropylium ring. Electronic Spectra. The electronic spectra of the ppolyphenyltropylium carbonium ions in four solvents are summarized in Table 11. As in the p-polyphenyl series,17 the successive addition of p-phenyl groups to the tropylium system produces a bathochromic shift of the longest wavelength absorption band. A straight-line correlation is noted between the observed frequencies for the longwavelength maxima of several tropylium salts obtained by us and other investigator^,^ and Am (the HMO calculated energy difference (in units of P) between the highest occupied and lowest unoccupied molecular orbital).lJ8 (See Table I11 and Figure 1.) Similar good correlations have been observed in other homologous conjugated systems.lg These bathochromic shifts can be attributed to an extension of the r-electron delocalization and the resultant decrease in the transition energy. Most importantly the long wavelength charge-transfer bands for 1, 2, and 3 show marked solvatochromic be-
2 3 4 5 6 11 12 13 14 15 16 17 18
p-C,H,-C,H,p-C,H,-C,H,-C,H,p-CH,O-C,H,Cyclopropyl Isopropyl 2-Naphthyl m-C,H,-C,H,H p-CH,-C,H,p-HOC,H,p-FC,H,p-ClC,H,p-BrC,H,-
cm- ’
Am
I , of ArH,f eV
27 250 1.2004b1C 9.245 24 150 1.0042b,C 8.27h 23 150 0.9100c 22 990 8.20 31 200 10.53’ 36 500 23 300 0.9464b 8.12 27 600 1.0560b 8.27h 3 6 4 0 0 1.6920 25 450 8.82 22 990 8.50 26 850 9.19 26 740 9.07 26 420 8.98
a Sovent: acetonitrile. Reference 18. Reference 1. Reference 5. e W. von E. Doering and L. H. Knox, J. A m , Chem. SOC.,76, 3203 (1954). f K. Watanabe, J. Chem. Phys., 26, 542 (1957). g Compounds 4-11, ref 5. K. Watanabe and T. Nakayamo, ASTIA Report ADH. Basch, M. B. Robin, N. A. Keubler, C. 152934. Baker, and D. W. Turner, J. Chem. Phys., 51, 52 (1969).
havior, the magnitude of the bathochromic shift increasing with decreasing solvent polarity. Although a plot of Kosower’s 2 values,20a measure of solvent polarity, vs. transition energy of these maxima in three solvents shows some deviation from a straight line relationship, this plot does indicate a correlation between the magnitude of shift and solvent polarity. The solvatochromicbehavior of the p-polyphenylyltropylium ions can be interpreted in the same manner as other systems such as the pyridinium cyclopentadienylium21,z2 and pyridinium phenolates23 which also exhibit intramolecular charge-transfer transitions. The magnitude of this solvent shift is also a function of the electron donor capabilities of the substituent. Thus the shift in the long wavelength absorption bands of 1,2, and 3 when solvent is changed from CHBCN to CH2C12is 20,35, and 51 nm, respectively, the extent of electron delocalization in aryl substituents increasing in the same direction. The presence of electron-donating heteroatomic substituents also cause increased shift between solvents. For example, the p-methoxyphenylThe Journal of Physical Chemistry, Vol. 81, No. 1, 1977
66
M. A. Battiste, M. W. Couch, and R. Rehberg
tropylium cation (4) (see Table 111) exhibits a shift (31nm) which is comparable to that observed for 2. A comparison of solvent effects of the p-polyphenylyltropylium salts with the p-polyphenyl series indicates little similarity. The solvent effects in the latter system are small and hypsochromic as contrasted with the large, bathochromic shifts in the tropylium ions, a fact consistent with the typical a a*transition characteristics for ppolyphenyls. Although we have so far discussed the electronic spectra of the monoaryltropylium ions in terms of molecular orbital theory, these spectra can also be described either in terms of an electron transfer from the phenyl ring, a .rr-electron donor, to the seven-membered tropylium ring, a a-electron acceptor (7), or in terms of valence bond structure 8. The former viewpoint is more formalized;
-
140
QOQO 7
where I,, is the ionization potential of the donor, E, is the electron affinity of the acceptor, and C is the Coulombic electrostatic energy.31 Since, in the monoaryltropylium salts, E , remains constant and C can also be considered invariant, the frequency of the charge-transfer band (v) should be proportional to the ionization potential of the donor. A plot of the frequency of the long-wavelength maxima observed for a variety of monoaryl tropylium salts vs. the ionization potential of the corresponding aromatic hydrocarbons gave a linear correlation. (See Table I11 and Figure 2.) The deviations from this straight-line relationship can be readily explained, In the m-biphenylyltropyliumsalts, delocalization of the positive charge into the remote phenyl ring gives rise to a high energy structure and, as a result, the m-biphenyl group has, effectively, a lower ionization potential than does p-biphenyl and behaves more like an unsubstituted phenyl group. Phenol, as has been demonstrated in other studies, acts as a better electron donor than is indicated by its ionization potential.32 The results analyzed above for the monoaryltropylium ions clearly establish the tropylium ion as an effective and sensitive probe of the electron donor properties of substituents in direct conjugation with the seven-membered ring, Since the cyclopropyl ring is known to be an efficient electron donor in the ground state, it was reasoned that the tropylium ion might prove to be superior to either a The Journal of Phvsical Chernistrv. Vol. 81.
/
8
however, either description clearly depicts the essential charge-transfer nature of the transition. We can, in fact, consider 1,2, and 3 to be typical type I11 tropylium ions, their long wavelength absorption bands being attributed to substituent-to-ring charge-transfer e~citation.~ The breadth of the long-wavelength absorption bands of 1-3 also suggests a charge-transfer transition since such bands are characteristically broad.23 The concept of the tropylium cation acting as a aelectron acceptor is not without precedent. For example, intermolecular charge-transfer complexes have been shown to arise from interactions between the tropylium ion and various aromatic hydrocarbon^.^^-^^ Charge-transfer complexation between the seven-membered ring and the halide ion also occurs in solid tropylium halides and their methylene chloride solutions to give dark colors.2s Similar interactions are also postulated to occur in the hexa- and heptaphenyltropylium halide^.^^^^^ Mathematically, the energy of a charge-transfer transition can be given to the first approximation by the expression E O = hv = I p - E,- C
No. 1. 1977
.-i /
80
1
, ' 23
I
I
I
24
25
26
I
27
28
,
29
I
30
I
31
(crn-7
Yx
Figure 2. Observed frequencies (urnax)and ionization potentials (4) for monoaryl tropylium ions.
benzenoid or another carbonium ion probe of the electron-donating ability of the cyclopropyl group in the excited state. Although a cation, the tropylium ion should behave more like a phenyl group than a typical carbonium ion. Previous spectroscopic studies of phenylcyclopropanes and some related rigid model systems have established conjugative interaction between the phenyl and cyclopropyl ring^,^^,^^ however, the steric relationship between the two rings was originally considered to have little spectroscopic i m p ~ r t a n c e . ~ ~ More recent ultraviolet spectral studies of cyclopropyl nitroaromatics by Hahn and his c o - ~ o r k e r have s ~ ~ shown that the latter conclusion is erroneous and probably reflech the minimal conjugative response of a cyclopropane ring to the weak electron demand of aromatic hydrocarbons. The electronic spectral characteristics of neutral aromatic systems and the generally small (5-10 nm) bathochromic shifts induced by cyclopropyl substitution greatly hampers the further detailed study of geometric and substituent influences on cyclopropyl conjugation in the first excited state. This problem could be partly overcome by use of a positively charged chromophore, such as a carbonium ion, which elicits a much greater electronic response from the cyclopropyl ring. Thus, as previously reported by Den0 and c o - w ~ r k e r scyclopropyl ,~~ substitution at the terminal carbon of an allylic cation such as 9 results in a significant bathochromic shift of A,, when compared to the methyl substituted ion 10. There are, however, major disadvantagesin performing spectroscopic studies on such reactive cationic species as 9 which are largely negated in stable tropylium ion salts such as 5 and 6. 10
9 A m a x 309nrnic28.000)
hmox275nm I~11,0001
Ahlnm) 34
As revealed by the comparison of ions 5 and 6, in Table 111, cyclopropyl substitutions of the tropylium ion produced a significantly greater bathochromic shift of the long-wavelength maxima (AXMeCN53 nm) than did the
67
Infrared Spectra of TI+N03- Ion Pairs
identical substitutions in allylic ion 9. Furthermore the for 5 is essentially an order of magnitude shift in, , ,A greater than that observed for cyclopropyl aromatic hydrocarbons. By Harmon’s designation cation 5 would be classified as a type I1 substituted tropylium cation displaying moderate intramolecular charge-transfer character. The somewhat greater sensitivity of, , A of 5 to a change in solvent polarity (CH3CN to CH2C12)than the isopropyl ion 6 is consistent with this interpretation. Given the above results it would now appear that the tropylium cation should be the preferred chromophoric probe of geometric and substituent influences on cyclopropyl conjugations in the excited state. Photochemical reactions of these ions should also be of interest for future exploration.
Acknowledgment. Financial support of this research by the National Science Foundation and the Air Force Office of Scientific Research is gratefully acknowledged. References and Notes J. B. Williams, Ph.D. Dissertation, University of Florida, 1967.
G. Hohlneicher, R. Kiessling, H. C. Jutz, and P. A. Straub, Ber. Bensenges. Phys. Chem., 70, 60 (1966). M. W. Couch, M.S. Thesis, University of Florida, 1966. K. M. Harmon, “Carbonium Ions”, Vol. 4, G. A. Olah and P. von R. Schleyer, Ed., Wiley-Interscience, New York, N.Y., 1973, and references therein. C. Jutz and F. Voithenleitner, Chem. Ber., 97, 29 (1964). H. J. Dauben, L. R. Honnen, and K. M. Harmon, J. Org. Chem., 25, 1442 (1960). A. Cairncross, Ph.D. Dissertation, Yale University, 1963. A. P. terBorg and H. Kloosterziel, Recl. Trav. Chim., PaysBas, 82, 741 (1963). A. P. terBorg, H. Kloosterziel, and N. Van Meurs, Recl. Trav. Chim., PaysBas, 82, 717 (1963).
(10) A. P. terBorg and H. Kloosterziel, Recl. Trav. Chim., PaysBas, 84, 241 (1965). (11) H. Gilman and E. A. Weipert, J. Org. Chem., 22, 446 (1957). (12) N. L. Bauld, J. D. McDermed, C. E. Hudson, Y. S. Rim, J. Zoeller, Jr., R. D. Gordon, and J. S.Hyde, J. Am Chem Sm.,91, 6666 (1969); these authors also report the preparation of cyclopropyltropylium fluoroborate (5, X = BF,-) and NMR spectrum (D,O) but do not comment further on the spectral properties of the cation. (13) M. Hanack and H. Eggensperger, Annales, 863, 31 (1963). (14) K. Conrow, J. Am. Chem. Soc., 83, 2343 (1961). (15) D. W. A. Sharp and N. Sheppard, J. Chem. Soc., 674 (1957). (16) G. Fraenkel, R. E. Carter, A. McLachlan, and J. H. Richards, J. Am. Chem. Soc., 82, 5846 (1960). (17) A. E. Gillam and D. H. Hey, J. Chem. SOC., 1170 (1939). (18) G. V. Boyd and N. Singer, Tetrahedron, 22, 547 (1966). (19) A. StreRweiser, Jr., “Molecular Orbital Theory for Organic Chemists”, Wiley, New York, N.Y., 1961. (20) E. M. Kosower, J. Am. Chem. Soc., 80, 3253 (1958). (21) D. Lloyd and J. S.Sneezum, Tetrahedron, 3, 334 (1958). (22) E. M. Kosower and P. E. Klinedenst, Jr., J. Am Chem. Soc., 78, 3493 (1956). (23) G. Briegleb, “Elektronen-Donator-Acceptor-Komplexe”, SpringerVerlag, Berlin, 1961. (24) M. Feldman and S.Winstein, J. Am. Chem. Soc., 83, 3338 (1961). (25) M. Feldman and S.Winstein, Tetrahedron Lett, 853 (1962). (26) M. Feldman and S.Winstein, Theor. Chem. Acta, 10, 86 (1968). (27) H. J. Dauben, Jr., and J. D. Wilson, Chem. Commun., 1629 (1968). (28) K. Harmon, F. E. Cummings, D. A. Davis, and D. J. Diestler, J. Am. Chem. SOC.,84, 3349 (1962). (29) M. A. Battiste, J. Am. Chem. Soc., 83, 4101 (1961). (30) T. Barton, Ph.D. Dissertation, University of Florida, 1967. (31) J. N. Murrell, Quart R. Chem. SOC.(London), 15, 191 (1961). (32) H. H. Jaffee and N. Orchen, “Theories and Applications of Ultraviolet Spectroscopy”, Wiley, New York, N.Y., 1962. (33) M. T. Rogers, J. Am. Chem. Soc., 69, 2544 (1947). (34) A. L. Goodman and R. H. Eastman, J. Am. Chem. Soc., 86, 908 (1964). (35) R. C. Hahn, P. H. Howard, S. M. Kong, G. A. Lorenzo, and N. L. Miller, J. Am. Chem. SOC.,91, 3558 (1969); R. C. Hahn, P. H. Howard, and G. A. Lorenzo, J. Am. Chem. Soc., 93, 5816 (1971). (36) N. C. Deno, H. G. Richey, Jr., J. S.Liu, D. N. Lincoln, and J. 0. Turner, J. Am. Chem. Soc., 87, 4533 (1965).
Infrared Spectra of TI+NO,- Ion Pairs Variably Hydrated or Ammoniated in an Argon Matrix G. Ritrhaupt and J. P. Devlln” Department of Chemistry, Oklahoma State University, Stillwater, Oklahoma 74074 (Received May 20, 1976; Revised Manuscript Received November 8, 1976) Publication costs assisted by the National Science Foundation
The vapor phase ion pairs, TI+N03-,have been isolated in argon matrices containing varying amounts of water or ammonia. The infrared spectra for these systems show quite clearly the effect of stepwise coordination of the T1+ ion by the solvent molecules. In particular, the splitting of the degenerate v3(e)nitrate mode, which is 275 cm-l for a pure argon matrix, is reduced in rather obvious steps to the limiting values of 53 and 18 cm-’ for the contact ion pairs in pure glassy H 2 0 and NH3, respectively. The manner in which this splitting collapses is markedly different from the smooth reduction reported previously for Li+N03-and this contrastingbehavior is analyzed in terms of the more covalent character of the bonding of T1+to NO3-, together with an apparent tendency for the Tl’ to move to the “roll on” position as it becomes solvated.
Introduction Several recent papers have emphasized that the volatility of polyatomic anion salts of the alkali metals, in particular the nitrates, chlorates, and perchlorates, makes it possible to position the corresponding ion pairs in a great range of environments for spectroscopic study.’ In particular data have been reported for ion pairs isolated in argon matrices and variably solvated in argon matrices containing water or ammonia.lc Like the alkali metal nitrates, thallium nitrate is known to volatilize associatively
so the vapor phase ion pairs, Tl+N03-, are readily condensed for matrix isolation spectroscopicmeasurements.la~d Thus, the ion pair vibrational spectrum has been reported for pure argon matrices with the most notable feature being the unusually strong distortion of the NO3- ion as reflected in the magnitude of the v3(e)doublet splitting.’* This splitting (275 cm-’) indicated that, although T1+ resembles the K+ ion in size, the thallium cation is as effective as the much smaller Li+ ion in distorting the nitrate anion and, further, that the Tl+--N03-interaction The Journal of Physlcal Chemistty VoL 81, No. 1, 1977
