NOTES
506
I-oi. G i MOLECULAR CORIPLEXES OF NETHOXYBESZESES BYARXOLD ZWEIG
Chemacal Research Department Central Research Dzurszon, Amerzcan Cyanamzd Company Stamford Conn.
100 -
Receawd August 6 , 1962
80
-
9
60-
d
-
0
Fig. 3.-The
I
CH,CN
0.2 0.4 0.6 Volume traction of polar component.
0.8
1.0
mutual solubility of CiFljH and polar liquids.
the last column, are the solubility parameters of the hydrocarbons and polar liquids a t 2.5’ based on the heats of vaporization and taken from Hildebrand and Scott.6 (The values for decane, dodecane, and tetradecane were interpolated by means of the linear plot of 61” us. the reciprocal carbon number of the paraffins n, the limiting value for l/n = 0 being 8.45.) Inspection of the results shows that the behavior of C7F16H in paraffins and alkylbenzenes does not differ markedly from that of perfluorocarbons. The paraffin solutions display the usual fluorocarbon “anomaly,” 61 being appreciably larger than &*. In alkylbenzene solutions the values of 6lagree quite well with the values of al*, which is also true in alkylbenzene-perfluorocarbon mixtures. However, in solutions containing a polar component marked differences become apparent. The critical solution temperatures of the systems C7FI5H-CH3KO2and C7F15-CH3CN are surprisingly low considering the solubility parameters of the liquids. Moreover, further experiments with polar solvents have shown that C7F15H-in contrast to perfluoropropylpyran-is completely miscible a t room temperature with methanol, ethanol, and acetone. On the basis of these results we may conclude that the behavior of C7FlBHin paraffins and alkylbenzeiies is dominated by the fluorocarbon chain. In polar solvents, however, specific interactions involving the highly polarized CH group of C7FI5Hstrongly influence or even dominate solution behavior. Acknowledgments.-The authors are indebted to Dr. C. A. Reilly for the nuclear magnetic resonance measurements. (6) J. H . Hildebrand and R L. Scott, “The Solubiiity of Non-Electrolytes,” Reinhold Publ. Corp., Neir I‘ork, S . Y., 1955.
Recent success in correlation of the orientation effects in the Birch reduction of methoxybenzenes with molecular orbital theory’ has suggested the examination of some other possible correlations of properties of methoxybenzenes with the simple Huckel molecular orbital (HAIO) scheme.2 AIolecular r-complexes of organic compounds have been known since 1858,3although their spectral propertied were not fully appreciated until the work ot Benesi and Hildebrand.4 hIulliken6 worked out a detailed quantum-mechanical theory for complexes where one molecule acts as an electron donor and another as an acceptor. In this theory, the charge-transfer absorption is ascribed to a transition from a ground state which is mostly an uncharged aggregate, DA, to an excited state which is mostly an ion-pair aggregate D+-A-. Dewar and Lepley6have shown that if the interactions between donor and acceptor are weak, that is, the ground state has little charge-transfer character, the transition energy, AEo, for the first charge transfer band for a series of donors with a given acceptor should be a linear function of the energies of the highest filled molecular orbitals (HFAIO) of the donors, assuming that the ionization potential is equal to the energy of the HFMO. Similar results with ionization potentials had been reached by Bhattacharya and Basu.’ Examination of charge-transfer absorption of catacondensed hydrocarbons with i ~ d i n e ,1,3,5-trinitro~ benzene,6 and tetracyanoethylenes have confirmed the linear relationship of AEo with the ionization potential and the HFMO of the hydrocarbon donor. Results and Discussion The major results of this investigation are summarized in Table I. The energies of the HMO’s mere obtained for each of the methoxybenzenes by assuming a planar delocalized system that includes six electrons from the benzene ring and two electrons from each attached oxygen atom. The carbon-oxygen exchange integral was taken as O.8pocand the oxygen 2pcc.2 coulomb integral was taken as CY, The secular determinant was factored when possible and the eigenvalues of the residual determinant were determined by an iterative procedure on a Burroughs 205 computer. Plotting the frequency of the chargetransfer absorption maximum against the HFhIO, as shown in Fig. 1, gave a very good correlation for seven of the methoxybenzenes and scattered points for the five others. From the slope, the indicated value of p is - 3 . 3 e.v. or -76 kcal./mole. Extinction coefficients and equilibrium constants
+
(1) H E. Zimmerman, Tetrahedron, 16, 169 (1961). (2) A . Streitwieser, “iUolecular Orbital Theory,” John J T h y & Sons, Inc , New York, 3’ Y., 1961. (3) J. V. Fritsche. J. prakt Chem , 73, [ l ] 282 (1858). (4) H A. Benesi and J H. Hildebrand, J . Am Chem Soc , 71, 2703 (1949). (j) R S Mulliken, z b z d , 74, 811 (1952). (6) AI J S Dewar a n d A. R Lepley zbzd , 83, 4560 (1961) (7) R Bhattachar,a and S Basu T r a n s Faraday Soc , 64, 1286 (1‘358) ( 8 ) 11 J. S D e u a r and H Rogers, J A m Chem S o c , 84, 393 (1962)
NOTES
Feb., 1963
507
TABLE I CHARGE-TRANSFER BAXDSOF COMPLEXES FORMED FROM METHOXYBENZENES AND TCNE BND THE HMO-HFMO VALUESOF THE METHOXYBENZEKES WaFe length of C-T band, mp
Compound
Frequeney of C-T band, Energy of ern.-' HFMO," X 10-8 xi8
Wave length of second C-T b a n d , mp
Methoxybenzene (anisole) 50gh 19.7 -0 827 387 1,2-Dimethoxybenzene (veratrole) 592 16.9 - 726 433 1,3-Dimethoxybenzene 548 18.2 - 773 470 (sh) - 692 380 1,4-Dimethoxybenzene 621 16.1 1,2,3-Trimethoxybenzene 515 19.4 - 705 1,2,4-Trimethoxybenzene 686 1 4 . 6 - .637 447 1,3,5-Trimethoxybenzene 552 18.2 - . 7 i 3 1,2,3,4-Tetramethoxybenzene 560" 1 7 . 9 - .627 480 l12,3,5-Tetramethoxybenzene 605 16.5 - .611 l12,4,5-Tetramethoxybenzene 800 1 2 . 5 - ,557 440 610 16.4 - .557 Pentamethoxybenzene Hexamethoxybenzene 512 19.5 - ,557 a The energy of the ith orbital is given by a zip. The i t h orbital here is the ( n 3) orbital whore n is the number of methoxy substituents. * For comparison, ref. 9 reported this band a t 507 mp. c The charge-transfer absorption was a broad plateau extending from 350 to 800 mp.
+
+
of the charge-transfer complexes by the Benesi-Hildebrand method4were not measured. Attempts to make such measurements on the hexamethoxybenzene complex were unsuccessful due to the weak absorption of the complex and the moderate solubility of tetracyanoethylene in methylene chloride. When equimolar concentrations of tetracyanoethylene and hexamethoxybenzene in methylene chloride (a pale purple solution) were slowly concentrated, the precipitate that, formed was pure tetracyanoethylene. On the other hand, the deeply colored 1,2,4,5-tetamethoxybenzene-telracyanoethylene solution in methylene chloride gave on concentration and addition of hexane, fine blue needles, m.p. 113-115', whose analysis indicated a 1 : I donor: acceptor ratio. A compressed pellet of this complex had a resistivity of 3 X 1O1O ohm.-cm. Examination of Fig. 1 shows that the complexes of all five compounds with three or more adjacent methoxy groups did not fit a correlation line with the calculated HFMO values while all seven compounds, ranging from monosubstituted to tetrasubstituted, with less than three adjacent methoxy groups fit the correlation line quite well. The only methoxybenzene for which crystallographic data are available, p-dimethoxybenzene, is reportedIc to have both methyl groups situated in the plane of the ring. It has been pointed out, from kinetic data and from dipole moment measurements, that the extent of conjugation of a methoxy group with the benzene ring is much reduced by the presence of two flanking Examination of models suggests that if methoxy groups are situated on three (9) R. E . Merrifield a n d W. D. Phillips, J A m Chem Soc., 80, 2778 (1958). (10) A. K. Kitaigorodskii, "Organic Chemical Crystallography,'' Translated b y Consultants Buieau. Nen P o r k , N. Y., 1961, p. 357; Acta Cryst., 3, 279 (1950). (11) H. P . Crocker a n d B. Jones, J. Chem. Soc., 1808 (1969). (12) K. B. Everaid a n d L. E Sutton, zbzd., 16 (1951). (13) W. F. Ansillotti and B. C. Curran. J . Am. Chem. Soc., 65, 607 (1943).
"t ,O6
0-0-
i
"0:
i
"4 I -1 5.0
14
I5
I
16 FREOUENW OF CHARGE-TRANSFER
1
I
17
18
BWD (an:' x
I 19
I 20
10-3).
Fig. 1.-Plot of the transition energies of the charge-transfer band of the methoxybenzenes against the energies of their highest occupied molecular orbitals.
adjacent sites in a benzene ring, the central methyl cannot lie in the benzene plane. With the methyl group out of the plane of the benzene ring, the nonbonding electron pair on oxygen is no longer able to participate to the same extent in the delocalized benzene ring system and thus the HMO energy levels, as calculated, are invalid. It thus appears that the HMO method is capable of predicting eigenvalue as well as eigenvector properties of methoxy substituted aromatics as long as steric requirements are met. E~perirnental~~ Materials.-Anisole was obtained from the Fischer Chemical Co. The three dimethoxybenzenes were obtained from Eastman Kodak Co. (White Label), and 1,2,4-trimethoxybenzene was obtained from Aldrich Chemical Co. These materials were used as supplied without further purification. Phloroglucinol (1,3,5)-trimethyl ether, m.p. .52-53', reported15 m.p. 52-53', was obtained by alkaline methylation of phloroglucinol,l5 and pyrogallol (1,2,3)-trimethyl ether, m.p. 45-46', reported16 m.p. 47", was obtained by a similar procedure.16 1,2,3,4-Tetramethoxybenzenewas obtained by partial methylation of gallacetophenone to 2-hydroxy-3,4-dimethoxyacetophenone followed by Dakin oxidation and further alkaline methylation, according to a published procedure,17 to the desired product, m.p. 88-90', reported17 m.p. 89.5-90'. 1,2,4,5-Tetramethoxybenzene was obtained by treatment of 2,5-dihydroxybenzoquinone (Eastman Chemical Products, Inc.) in anhydrous methanol with hydrogen chloride t o give 2,5-dimethoxybenzoquinone,~* which was then reduced and treated with alkaline methyl sulfate,'? m.p. 102-103', reported" m.p. 102-103". 1,2,3,5Tetrametboxybenzene was prepared by the method of Baker18 through the intermediacy of pyrogallol trimethyl ether and 2,6dimethoxyquinone, m.p. 40-42', reported19 m.p. 47". Pentamethoxybenzene was prepared from 1,2,3,5-tetramethoxybenzene by Baker's method,'$ m.p. 55-57', reportedIg m.p. 58-59'. Hexamethoxybenzene was prepared from 2,6-dimethoxybenzoquinone by the method of Robinson and Vasey,20 m p . 81', reportedz0m.p. 81". Thus, five of the seven prepared methoxybenzenes were made from pyrogallol. Yields were comparable t o those reported in the literature. Infrared and ultraviolet spectral examination of all twelve methoxybenzenes were consistent with the proposed structures and did not indicate any obvious contaminations. Tetracyanoethylene was purified by recrystallization from chlorobenzene and sublimation.9 The methylene chloride used was spectral grade. (14) All melting points are uncorrected. (15) J. W. Clark-Lewis, Az~stralianJ. Chem., 10, 505 (1957). (16) E. Chapman, A. G. Perkin, a n d R. Robinson, J. Chem. Soc., 3028 (1927). (17) F. Benington, R. D. Morin, a n d L. C. Clark, Jr., J. Or#. Chem., 20, 103 (1955). (18) R. Scholl a n d P. 13ahl1, Ber.. 67, 83 (1924); A. H. Crosby a n d R. E . L u t z , J. Am. Chem. Soc., 1 8 , 1235 (1956). (19) U ' . Baker, 6.Chem. See., 662 (1941). (20) R. Robinson a n d C. Vasey, %bad.,660 (1941).
508
NOTES
Vol. 67
Spectroscopic Measurements.-Concentrated solutions of the rnethoxybenzenes and tetracyanoethylene in methylene chloride were mixed in approximately equimolar proportions, then diluted, if necessary, to obtain the position of the absorption maxima. Measurements usually were made in a C a y Model 11 specbrophotometer for the 300-800-mp range. The tetra-, penta-, and hexamethoxybenzene complexes also were examined in the 700-1500-mp range on a Cary 14 spectrophotometer.
less the energies of breaking Cu-0 bonds and Si-Xi bonds are large. Fortunately, it has been shown that ethylmethylglyoxime forms a complex with nickel(II), [Ki(EMG)z], which is similar in every way to n'i(DRIG)2 except that in Ki(EMG)2 there are no Xi-Xi bonds, presumably because of steric fact01-s.~Hence the enthalpy of soluAcknowledgment.-Thanks are due to J. H. Lehnsen tion of Ki(EhfG)zin heptane should be about equal to for assistance in the computations and to the Ultrathat of Si(DMG)2if there were no Ni-Ni bonds in the violet Spectroscopy Croup of the Research Service latter. This value for Ki(ERIIG)z,which should also be Department of the Stamford Laboratories for their a good estimate of AHo of solution of C U ( D M G ) ~ cooperation. monomer, 5.8 kcal./molej6 leads to an estimate of the energy of the Si-Si bonds of -9 kcal./mole. (Values for solution in CC14 lead to an estimate of about 11 Or\; SOLUBILITIES AND STRUCTURES kcal./mole.)6 OF S I C K E L AND COPPER Fleischer and Freiser have suggested that the Xi-Xi DIMETHYLGLYOXIMES bond distance, 3.25 A., far longer than the estimated single bond distance, 2.31 A,, is in disagreement with BY R. E. RUNDLE A N D C. V. BAXKS any appreciable n'i-Ni bonding, but long metal-metal Contizbulzon 1182 h m the Instztute .for Atomic Research and Department bonds are known elsewhere. In lLInz(CO)lothe Mn-hln of Chemzstry, Ioua State Unzverszty, Amee, Ioua distance is 2.93 A.7 us. 2.35 A. for an estimate of the Recewed June 21, 1962 single bond distance, yet the bond energy has been Fleischer and Freiser2 recently have determined the found to be 34 f 13 kcal./mole.8 solubilities and heats of solution of nickel and copper I n any discussion of the solubility data, the effect dimethylglyoximes, [i"\'i(D,lilG)zand Cu(Dh1G)z 1, in a of Cu-0 bonds in crystalline CU(DR!IG)~ and the possivariety of solvents, and have interpreted their data to bility of a large change in hydrogen bonding in going mean that Xi-Ni bonds in Ki(DT\fG)zare far weaker ~ monomers in heptane than the previously estimated value of -10 k ~ a l . / m o l e . ~ from crystalline C U ( D M G ) to solution must be taken into account. Each of these The recent structure determination which shows Cufactors should be large and in opposite directions. (DR1G)z to consist of dimers in the crystal, bonded toSince both will be uncertain, all that can be concluded gether by Cu-0 bonds,4 raises many complicating with fair certainty is that the difference between these features in any interpretation of the solilbility data. quantities must be about 11 kcal./mole to account for We show here that these data are consistent with -10 the large, 16.6 kcal./mole, enthalpy of solution of Cukcal./mole for Ni-Ni bonds if proper recognition of the (DT\!IG)zin heptane. We estimate bond energies below crystal structures of the two compounds is made. mainly to point out the complications and further inFurthermore, in considering the structure of Cu(D;LfG)z, terest in this problem. some very interesting points arise which deserve further The hydrogen bonding in crystalline Cu(DMG)* is study. very different from that in S(DMG)2. In the crystal, The previously reported entropies of solution were Fig. 1, each monomer of Cu(DMG)2 has two hydrogen obtained by dividing AHo by T , as though the Gibbs bonds of 2.70 and 2.53 A.,4while there are two hydrogen free energy of solution, AGO, were zero,2 which is far bonds of 2.40 A. in Ni(DMG)z.s In crystalline Cufrom the case for these very slightly soluble compounds. (DhIG):! each oxygen, 011,is bonded to Cu of the neighCorrected thermodynamic values are given in Table I . boring monomer to form dimers with two Cu-0 bonds The data of Table I suggest that Si(DRIG)2 and per dimer. Undoubtedly the Cu-0 bonds lead to the C U ( D M G )are ~ essentially organic type solutes. Neverweaker and longer 0-H-0 bonds in Cu(DNIG)2 crystals, theless, the enthalpies and entropies of solution suggest but in Cu(DMG)2 monomers in solution, strong hydrothat both Ni(DMG), and CU(DNIG)~ form hydrates gen bonds of the type found in Ki(DMG)2 are to be with water, lowering the enthalpies of solution in water, expected. and tying up water so as to lower the entropy of solution. The enthalpy involved in this rearrangement of These effects of hydration are much more important hydrogen bonds in C U ( D ~ I Gis)likely ~ to be very large. for CU(DT\IG)~than for Ni(DMG)2 as noted by The 0-H-0 bonds of 2.40 A. are probably symmetrical Fleischer and Freiser. The conclusion that similar hydrogen bonds, because neutron diffraction has shown solvation effects are very important in chloroform and that other hydrogen bonds of about this length are are still prominent in benzene also seems inescapable. symmetrical.10 Hence, in energy these 0-H-0 bonds Like Fleischer and Freiser, we suppose that the solvaare to be compare$ with the symmetrical F-H-F hydrotion effects must be negligible in heptane, and that in gen bonds of 2.27 A. found in KHFZ. For this hydrogen heptane the molecules of Si(DhlG)zand Cu(DMG)*are bond, the bond energy has been estimated to be about nearly identical. It then appears that the enthalpies of solution of Ki(DMG)z and CLI(DMG)Zare very large (5) E. Frasson and C. Panattoni, ibid., 13, 893 (1Q60). (6) C. V. Banks and S. Anderson, J . Am. Chem. Sac., 84, 1486 (1962). (14.8 and 16.G kcal./mole, respectively) in heptane un(1) Work \\as performed in the ?mes Lahoratory of the U. S. Atomic Energy Commission. (2) D. Fleischer and H. Freiser, J P h y s . Chem , 6 6 , 389 (1962), see also the solubility data of D. Dyrssen and MI. Hennichs, Acta Chem Scand 15, 47 (1961). ( 3 ) L E. Godycki and R. E. Rundle, Acta C r y s t , 6, 487 (1953). (4) E. Frasson, R Bardi, and S. Bezai, z b d , 12, 201 (1959).
(7) L. F. Dahl, R. E. Rundle, and E. Ishishi, J. Chem. Phys., 26, 1750 (1957). (8) F. A. Cotton and R . R. illonohamp, J . Chem. Soc., 533 (1960). (9) D.E. Williams, G. Wohlauer, and R. E. Rundle, J . Am. Chem. SOC., 81,758 (1Y59). Levy, J. Chem. P h w , 29, 948 (1958); and (10) S. \T, Peterson and H Levy. as quoted in 1,. Pauling, "Nature of the Chemical Bond," 3rd E d , Cornell Univ.. Press, Ithaca, N. Y., 1958, p. 948.
