Oct., 1963
I N T E R A C T I O N B E T W E E K R f E T H Y L B E h ’ Z E S E S AND POLYCYCLIC H Y D R O C A R B O S S
220 1
MOLECTJLAR INTERACTION BETWEEN METHYLBENZENES AND POLYCYCLIC AROMATIC HYDROCARBONS BY
w.E.
W E N T W O R T H AKD ELWARD C H E S
Department of Chemistry, University of Houston, Houston Received April 10, 1963
4, Texas
The change in the spectra of the three- and four-ring polycyclic aromatic hydrocarbons upon addition of methylbenzenes in an inert solvent has been interpreted as due to complex formation. The observed increase in the complex stability with decreasing ionization potential of the methylbenzenes is considered as evidence for a t least some charge-transfer stabilization. The correlation of increased stability with increasing electron affinity of the polycyclic aromatic hydrocarbons adds further proof to the charge-transfer interpretation. In turn it also adds support to a previous study on a proposed experimental method for the determination of molecular electron affinities.
Introduction Experimental determinations of the electron affinity of molecules are very rare, but recently the experimental measurements of electron capture coefficients of several aromatic hydrocarbons have been interpreted by Wentworth and Beckerl in terms of the electron affinities of the molecule. In a, later paper2 these experimental electron affinities hare been correlated with the lowest energy 0-0 frequency transition of the hydrocarbon and a,lsowith theoretical calculations. Previously, the only experimental estimates of electron affinities of large organic molecules came from the interpretation of polarographic half-wa,ve reduction potentials which are complicated by the uncert’aiiities in the solvat’ioii energy. The ;11ullikei1~~~ theory of charge-transfer complexes predicts that the frequency of the charge-transfer band and the stability of the various complexes are related to the ionization potential of the donor and the electron affinit,y of the acceptor. d considerable amount of work6.6 has been done correlating the frequencies and the stabilities with the ionization potential of the donor for a series of complexes in which the acceptor is kept constant, but because of the scarcit,y of data on electron affinities of molecules, very little has been done with a series of acceptors and the correlations with their electron affinities. Briegleb and C ~ e k a l l a ~have - ~ assumed the theory to be correct and have actually obtained electron affinities from frequency, stability, and dipole moment dat’a for cha,rge-transfer complexes. Jortner and SokolovlO and Lyons and Batley” have made tlie same assumption and have used the electron affinity of the iodine atom and the frequency of the charge-transfer band for complexes in which iodine acts as the acceptor to establish a scale of relative electron affinities. P e o ~ e r ’ ~has . ’ ~ shown that the stability and the frequencies of Complexes of a series of quinones with a com(1) W. E. Went~vortliand R. S. Becker, J . Am. Cicem. Soc., 84, 4263 (lll62). (2) R. 8. Recker a n d W. E. Wentworth, ibid., 85, 2210 (1963). (3) R. S.Mulliken, ibid., 72, 600 (1950). (4) R. S. Mulliken, ibid., 74, 811 (1952). 15) J. L. Hastings, J. Franklin, J. Shiller, and F. A. Matsen, ibid., 75, 2900
(1953). ( 6 ) H. McConnell, J. S. Barn, and J. R. Platt, J. Chem. Phus., 21, 66 (1953).
(7) G. Briegleb, “Elecktroileii-Donator. Aocagtor Komplexe,” SpringerVerlag, Berlin, 1961. ( 8 ) G . Brievleb and J. Czekalla, Z. Elektrochem., 63, 1157 (1959). (9) G. Briegleb and J. Czekalla. Anffew.Chem., 72, 401 (1960). (10) J. Jortner and W. Sokolov, .\‘atwe, 190, 1003 (1981). (11) M. liatley and L. E. Lyons. ibid.. 196, 575 (1962). (12) i\I. E. Peover, ibid., 191, 702 (1961). (13) 1LI. E. Peover, Trans. Paradav Soc., 58, 1656 (1962).
mon donor can be correlated with the half-wave reduction potential of the quinones which in turn have been correlated with the electron affinities of the quinones. However, again, the uncertainty in the solvation energy prevents the correlation with absolute electron affinities. Therefore, there is still a need to investigate the correlations of charge-transfer frequencies and the thermodynamic properties with the independently determined electron affinities of the acceptor molecule. Since the experimental values for the electron affinities of the aromatic hydrocarbons are ~ O T Vavailable, it appeared desirable to study charge-transfer complexes in which the polycyclic aromatic hydrocarbon acts as the acceptor (1) in order to establish experimentally the validity of tlie theoretical relationships between the charge-transfer parameters and the electron affinity of the donor and ( 2 ) to support further the interpretation of the electron capture measurements in terms of the electron affinities of the molecules. Numerous charge-transfer complexes involving polycyclic aromatic hydrocarbons as donors have been reported in the literature. On the other hand, the evidence for complex formation in which polycyclic aromatic hydrocarbons act as electron acceptors is very limited. I n a recent paper spectral evidence was used for establishing a charge-transfer complex between acridine and carcinogenic compounds.14 I n this case the carcinogens, including various polycyclic aromatic hydrocarbons, were acting as electron acceptors. Khanna, et al., 15.16 have reported a charge-transfer complex in which “living” polystyrene, a polymer with a free electron, acts as a donor with anthracene as the acceptor. Thus, the first step in the present study was to establish the existence of complexes with these aromatic hydrocarbons acting as acceptors. The second step was the determination of the equilibrium constant and hence the change in standard free energy of the complex. The frequency of the charge-transfer band could not be investigated because a comparison of the electron affinities (see Table I) of the aromatic hydrocarbons with those of the common charge-transfer acceptors indicates that the bands would occur in the farultraviolet region where both the donor and acceptor molecules have intense absorption bands, making the observation of a new band very difficult. A study of a series of methylbenzene donors with one acceptor, an(14) A. C. Allison and T. Nash, Nature, 197, 758 (1963). (15) S. IChanna, M. Levy, and >I. Szwarc, Trans. Faraday Soc., 58, 747 (1962). (16) R. Asami, S. Khanna, M. Levy, and M. Szwarc, ibid., 68, 1827 (1962).
2202
W. E. WENTWORTH AND EDWARD CHES
thracene, was conducted in order to lend support to the charge-transfer contribution to the stability of the complex. The final step was the correlation of the thermodynamic data with the electron affinities of the acceptors. Experimental Chemicals.-The aromatic hydrocarbons used as acceptors were the same samples as those used by Becker, et al.,'' which were analyzed for purity by mass spectrometry and the results reported. Eastman's Vhite Label p-xylene, durene, and hexamethylbenzene and Fisher's Certified Reagent toluene were used as donors without further purification. Only one major peak was observed when the purity of the p-xylene and toluene was checked by gas-liquid chromatography. From the gas chromatographic results, the concentrations of any suspected impurities must be a t such a low level that they could not seriously affect the results. In the case of durene and hexamethylbenzene, the effect of small impurities of substances of weaker complexing ability with the aromatic hydrocarbons would be negligible. Phillip's 99 mole yo heptane was passed through a silica gel column and spectroanalyzed in the ultravioiet before use as the solvent. Spectra.-Difference spectra for a series of solutions with a given concentration of acceptor and a varying concentration of donor were obtained using a solution of acceptor in heptane of the same concentration as that on the sample side as a blank. The acceptor concentrations were approximately M and the donor concentrations varied from 0.1 to 0.46 mole fraction unit except in the case of durene and hexamethylbenzene, where the concentration of donor was limited to 0.1 mole fraction unit by solubility. The solutions were prepared immediately before making the measurements. A Cahn electrobalance was used to weigh out on the order of 6 mg. of the aromatic hydrocarbons to a precision of mg. The liquid donor solutions were prepared volumetrically. The temperatures ranged from 27.5 to 32.0" from run to run, but remained essentially constant during a run. The qualitative spectral data were obtained on a Beckman DB recording spectrophotometer and/or a Bausch and Lomb Spectronic 505 recording spectrophotometer. In addition to the simple difference spectra, qualitative data were obtained for double difference spectra in which two cells were used in both the reference and the sample sides. One cell on the reference side contained a solution of the acceptor in heptane and the other contained a solution of donor in heptane. In the sample beam, one cell contained a solution of donor plus acceptor in heptane and the other contained pure heptane. This combination canceled out the absorption of the donor and the acceptor and ensured that the spectra being measured was due t o the mixture of acceptor and donor. The quantitative data used to determine the equilibrium constants were obtained on a Beckman DU spectrophotometer. The cell compartment was not temperature controlled and the spectrophotometer was not calibrated for wave length. The wave lengths a t which the quantitative data were obtained were selected by obtaining the maximum difference in absorption between the mixture of polycyclic aromatic hydrocarbons plus 0.1 mole fraction of donor and the solution of identical concentration of aromatic hydrocarbons in heptane. This generally was in the region to the long wave length side of the longest wave length band of the polycyclic aromatic hydrocarbons. The slit widths and sensitivities were adjusted until the transmission remained constant with a decrease in slit width and an increase of sensitivity. The slit width was in all cases less than 0.01 mm. In the case of chrysene and benzo(c)phenanthrene, the shift of the longest wave length band was not used. The quantity of pure samples of these compounds was quite limited and in order to use smaller quantities of material, the shift in a more intense band a t shorter wave lengths was selected. I n this wave length region, p-xylene has a weak absorption and it was thus necessary t o obtain the extinction coefficient of p-xylene a t the wave length used. This was done on the DU by measuring the absorption of a series of solutions of p-xylene in heptane using heptane as a blank. Data t o determine the equilibrium constant for the benzeneanthracene complex were obtained, but because of the small (17) R. S. Beckel, I. 2141 (1963).
S. Hingli, and L. 1. Jackson, J.
Chem P h y a , 38,
magnitude of the equilibrium constant, i t was not possible to obtain meaningful results from the data reduction process. As will be discussed in the following section, the least squares procedure involves a Taylor series expansion. The calculation procedure consists of successive iterations until the estimates of the parameters converge. For extremely weak complexes, the parameters are not as m-ell defined and difficulties are experienced in obtaining satisfactory estimates t o make the process convergent.
Calculation of Equilibrium Constants As mill be discussed subsequently, the change in the absorption spectra for solutions of polycyclic aromatic hydrocarbons in mixtures of methylbenzenes in heptane has been interpreted in terms of molecular complex formation. The quantitative spectrophotometric data were reduced on an I B N 7090 digital computer using a program for the least square adjustment of experimental data in order to calculate the equilibrium constant for the complex formation. The spectrophotometric data were obtained from difference spectra as described in the Experimental section. These data, along with the known initial concentrations, were used as input and the equilibrium constant and two other parameters were obtained. The mathematics of the program is given in the following paragraphs. The data points are properly weighted based upon the assumption that the variance in the transmittance readings is constant, independent of the magnitude of the transmittance. Many investigators overlook this important feature of weighting in least squares in the calculation of the equilibrium constants. It is assumed that a one to one complex exists between the methylbenzenes and the polycyclic aromatic hydrocarbons. Representing the polycyclic aromatic hydrocarbon by A and the methylbenzene by B, the complex formation thus is A+B=AB Assuming the activity coefficients are nearly unity, the equilibrium constant, K , expressed in terms of concentration is given by
For weak complexes where the concentration of B is in excess, the initial concentration is approximately the actual concentration of B in the solution. [B] a b = initial concentration of methylbenzene Letting a = initial concentration of polycyclic aromatic hydrocarbon, eq. 2 becomes
K =
[AB1 (a - [AB])b
(3)
Solving for the concentration of the complex [AB] =
~
1
Kab bK
+
(4)
Since the quantitative data were obtained from difference spectra, we must consider the absorbance of both the reference solution and the complex solution in relating the absorbance to the complex equilibrium expression, in eq. 4. The absorbance of the reference solution containing the polycyclic hydrocarbon (A) in heptane is given by
INTERACTIOX BETWEEN METHYLBENZENES AND POLYCYCLIC HYDROCARBONS
Oct., 1963
p r = LEAa
+
2203
(5)
Pro
where p r = absorbance of the reference solution, 1, = is the path length of the reference solution, EA = molar extinction coefficient of the polycyclic aromatic hydrocarbon, a = concentration of the polycyclic aromatic hydrocarbon in moleF/liter, and pro = absorbance of the reference cell. At the wave lengths selected, the methylbenzenes did not show any absorbance and therefore the absorbance of the complex solution is given by ps =
LEA1 9 1
+ LEAB[AB] +
(6)
pso
where the symbols used have the same connotation as above, except that they refer to the complex or sample solution. Expressing [A] in terms of the a and [AB], (6) can be rearranged to ps =
LEAU
-k
&(CAB
-
c-4)
[AB]
+
(7)
pso
The difference spectrum observed then is obtained by subtracting eq. 5 from eq. 7 AP = p s
-
Pr
=
(Is
- & ) e ~ a+ L(EAB- EA)
X
+ A P O (8)
where Apo = pso - pro. The path lengths of the cells were matched within. 0.001 cm. and the first term in eq. 8 can be neglected considering the magnitude of EA a t the selected wave lengths. Keglecting this term and substituting eq. 4 into eq. 8 we obtain
h,mu. Fig. 1.-Per cent T of 1 X lo-* M anthracene in heptane with varying concentrations of p-xylene expressed as mole fraction: (1) 0; (2) 0.10; (3) 0.30; (4)0.45; (5) 0.55; (6) 0.65.
This is the equation which was adjusted to the experimental data according to the principle of least squares.l* The parameters which were adjusted were (EAB - EA), K , and Apo. The condition equation in the least squares solution is
The weights of the ai and bi were assumed infinite (negligible error compared to the spectrophotometric measurements:) and the weights of Api are
where uo2 = variance of unit weight. Since lated to the transmittance by Api =
- log Ti
Api
is re-
(12)
where uT,2 = variance in transmittance = constant. From eq. 11, setting no2= (0.434)2 uT,2
W,,,,, = Ti2 ( 14) It can be seen that higher absorbance values result in lower Ti values and hence less weight in Ap,. This is in accordance with the greater absolute error in Apl as ApL increases. (18) IV. E. Demirig, "Statistical Adjustment of Expeniuentai I h t a , " Jolin Wile>- and Sons, Inc , New Yolk, N. Y . , 1943.
I ( '- _ 400 390 380 370 360 350 340 330 A, mu. Fig, 2.-Difference spectra of anthracene (C = 1 X M) in 70 T as a function of varying concentrations of p-xylene expressed as mole fraction: (I) 0; (2) 0.15; ( 3 )0.35; (4)0.45; (5) 0.55; (6) 0.65.
This least squares procedure differs from that employed by many other authors. I n their analysis, a least squares adjustment is made to a linear form of the Benesi-Hildebrand equation and generally the points
W. i3.WENTWORTH AND EDWARD CHEN
2204
100
DONORS
Mol. 67
IO0
ACCEPTORS
@ HEXAMETHYLBENZENE =ANTHRACENE fl DURENE
)PHENANTHRENE ANTHRACENE
dCHRYSENE
0 P-XYLENE
75
% T.
75
h-,
% T.
F"D3
5c
50
\
0"
25
\'..a -
I
.IO
I
I
.20
c
I
30 A0 (MOLE FRACTION).
0x1
I
.50 C
\
MOLE FRACTION ) .
Fig. 3.-Per cent T of difference spectra a t a fixed wave length as a function of concentration of donor. Solid line is least squares adjustment to the data.
Fig. 4.-Per cent T of difference spectra a t a fixed wave length as a function of concentration of donor, p-xylene. Solid line is least squares adjustment to the data.
are not properly weighted. One questions whether this is really a least squares adjustment a t all. As emphasized by Deming, the true least squares solution consists of a minimization of the sum of the weighted squares of the residuals of the observations.
equilibria, it would add two additional parameters to be adjusted. With this many parameters most certainly the data could be fit to these functions; however, it is questionable whether the experimental data are accurate enough to definitely establish whether this additional 2 : 1 complex is significant. For this reason, the authors have limited the treatment of the data to a 1: 1 complex. As expected, the charge-transfer bands could not be observed in this region of the spectra. The double difference spectra did not differ markedly from the single difference spectra so that the measured difference in absorption must be due to the combination of the donor and the acceptor. Perhaps the most convincing evidence for complex formation is the fit of the data to the expected equations for a 1:l complex. Figures 3 and 4 show graphs of per cent transmission vs. concentration of the donor. The approximate wave lengths a t which these measurements were made for each of the pairs of compounds can be found in Table I. The expected standard deviation in transmission, 0.005, is shown by the circles. The average standard deviation in transmission as computed by this program was 0.007 so that the fit of the data was to about the expected degree. Of course, this does not unambiguously establish the nature of the complex since, as B a y l i ~ s ' ~pointed out, the data probably would fit the equations for an n : 1 complex. The equilibrium constants and the uncertainties in these parameters expressed as standard deviations ob-
Results Figure 1 is an example of the spectra for a series of solutions in which the concentration of the electron acceptor, anthracene, is maintained constant and that of the donor, p-xylene, is varied. The reference solution is pure solvent. Figure 2 is the difference spectra for a series of solutions with an anthracene solution as the reference. Upon the addition of donor to acceptor solution, the spectra of all of the other aromatic hydrocarbons shifted as illustrated by the shifts of the spectra of anthracene, strongly indicating the formation of a complex. Other indications of complex formation are the isosbestic points which can be seen a t approximate wave lengths of 3750, 3710, 3665, 3.560, 3480, 3380, 3350, and 3370 A. These points are more easily seen in the difference spectra. The isosbestics may not be as well defined as one may expect for a 1: 1 complex. This possibly could be due to some formation of a 2 : 1 complex in which there are two methylbeiizenes associated with a single polycyclic aromatic hydrocarbon molecule. If the structure of the complex is such that the aromatic rings are planar, then with an excess of p-xylene the possibility of a 2 : 1 complex seems quite reasonable. However, if one includes this equilibrium of a 2 : 1 complex in addition to the 1: 1 complex in the least squares calculation of the
(19) N. 8 . Bayliss and C. J. Brackenridge, J . Am. Cham. Soc., 77, 3 Y X (1955).
act., 1963
INTERACTION BETWEEN S~ETHYLBFNZEKES AXD POLYCYCLIC HYDROCARBONS
2205
TABLE I CHANGE IN STAXDARD FREEENERGY ASD EQUILIBRIUM CONSTAKTS FOR AROMATIC HYDROCARBON COMPLEXES AF',
aAFo,
u~,inf.-l kcal./mole koal./mole Triphenylene p-Xylene 338 0.14 8.45 0.129 0.264 1.240 1.220 0.313 p-Xylene 348 .20 8.45 .098 1 096 Phenanthrene .164 0 408 .249 Chrysene p-Xy1ene 321 .33 8.45 ,509 .209 ,330 .240 320 .33 8.45 ,252 Benzo(c)phenankhrene p-Xylene .578 .704 .l54 Pyrene p-Xylene 8.45 .313 .080 346 .39 Anthracene p-Xylene .093 378 .42 8.45 .126 ,117 .825 Bene(a)anthracsne p-Xylene 387 .46 8.45 ,468 ,102 .460 .122 Anthracene Toluene 378 .42 8.82 ,122 ,055 1 27 .271 Anthracene p-Xylene 378 .42 8.45 ,825 .126 0.117 ,094 Anthracene Durene 378 .42 8.02 1 768 .631 -0 346 .215 Anthracene Hexamethylbenzene 378 .42 7.85 9.462 1.514 -1.36 .096 a Electron affinity measurements were taken from ref. 1. All ionization potentials have been reported by K. Watanabe [ J . Chem. Phys., 26, 542 (L957), and K. Watanabe, T. Nakayama, and J. Mottl, "Final Report on Ionization Potential of Molecules by a Photoionization Method," December, 1959, Dept. Army 5B99-01-004, ORD TB2-0001, OOR 1624, Contract No. DA-04-200-ORD 480 and 7371, except for hextimethylbenzene which was reported by Vilesov [F. I. Vilesov, Zh. Fiz. Khim., 35, 2010 (1961)j. Acceptor
Donor
X,mp
E A , ~ ~ . v . I ~ , ~ e . v . K,mf.-1
'
taiiied from the least squares program are shown in Table I. Molar extinction coefficients were also obtained, but no effort was made to select the wave length of maximum absorption so that they were not included. I n a consideration of the nature of the forces binding the complex, most certainly one must include the classical dispersioii or van der Waals forces. However, keeping in mind the experimental estimates of the eiectron affinities of the polycyclic aromatic hydrocarbons and the relatively low ionizatioii potentials of the methylbenzeries, it seems possible, at least, that chargetransfer forces coulld contribute significantly to the binding energy. Since no charge-transfer bands mere observed, the only correlation which one can make with these complexes is with the thermodynamic properties. According to Mulliken's3~*charge-transfer theory the resonance energy as a consequence of charge transfer is given by
where E N is the energy of the ground state referenced to the zero energy of the separated molecules; Eo is the energy of the no-bond state and includes electrostatic interactions such as (dispersionforces; E1 is the energy of the dative (charge-transfer) state before resonance interaction is considered ; HO1 is the resonance integral of the "no-bond" state and the dative state; S is the overlap integral of these same two states. The difference in energy (E, - Eo) has been expressed by Hastings, et aLj5as
where I B is the ionimtion potential of the base (methylbenzene), EA is the electron affinity of the polycyclic aromatic hydrocarbon, T A B is the separation of the charged molecules iii the charge-transfer state, and C-kU is the difference in th.e dispersion aiid dipole energies of the "no-bond" and dative states. Substitution of eq. 16 into ey. 15 yields
-1icGIyiin20has suggested tllak for a series of similar
complexes, Holis proportional t o S. If we furthermore assume that for a series of similar complexes X and (e2/rtlB CAB)are constant
+
Actually, X and HOl would not remain exactly constant; however, the variation in these quantities is probably small compared to the change in I n and Eg. Finally, for a condeiised system EN can be replaced by the change in enthalpy, AH", and for simple complexes as observed in charge-transfer21 or hydrogen-bondedZ2 complexes, AH" is found to be linearly related to AFO. We then hare an equation relating AF" with I B and EA which is AF"
=
c 1
I B - E A - Cz
+ ca
(19)
The constant Cs will also include the difference in solvation energies for the components aiid the complex. In a n inert solvent these should remain approximately constant. Over small ranges of I B and E.4, AF" should vary almost linearly with these parameters. This is especially true if the I B is relatively large and the EAis relatively small compared to the remaining terms in the denominator. As the I g becomes smaller, one would expect the stabilization to increase (AF" to become larger negatively). On the other hand, as the E A becomes larger oiie would expect the stabilization to increase. Briegleb7-9 actually has determined some of these constants for an analogous relationship of AH', but it was necessary to have both thermodynamic and chargetransfer frequency data. In cases where it is not possible to obtain frequency data, it is just as meaningful to examine AF" variations as it is to examine AH' dependencies, since AF" is most likely linearly related to AH". There may be some doubt as to whether these three constants C1, CZ,and CSremain truly constant. For a given set of similar compounds such as between the methylbenzenes and the polycyclic aromatic hydrocarbons, there is a good chance that they do not remain constant. The only real justification for these as(20) s. P. hIcGlynn, Che n. ReL., 68, 1113 (1958). (21) L. J. Andrews, zbzd , 64, 713 (1954). (22) G. C. Pimentel and A. L. hIcClellan, "The Hydrogen Bond," W. H. Freeman & Co., 1960.
ITT. E. WENTWORTH AND EDWARD CHEX
2206
Vol. 67 ACCEPTORS
& TRIPHENLYENE
2
PHENANTHRENE @' CHRYS ENE 88BENZO IC) PHENANTHRENE 8 PYRENE m, ANTHRACENE C& BENZ (n)ANTHRACENE DONORS HEX AMETHYLBENZENE nDURENE 0P- XYLENE GTOL UENE cCp
ACCEPT OR ANTHRACENE
2.0 -
-r-
HEXAMETWLBENZENE
1.0 + AF'
kca I.) ,
(m P -X Y L E IE
0. I
\
T
-LO--
I
I
8.50
8.80
L
9.00
I
I
I
I
.I8
.I9 I
.20
.21
(lg-EA-3,
Fig. 5.-Change in standard free energy for complexes of methylbenzenes with anthracene zis. the ionization potential of the methylbenzene.
DONOR P - XYLENE ACCEPTORS sb TRIPHENYLENE
.I5
.20
.25
.30
.35
.40
.45
E a (electron volts) Fig. 6.-Change in standard free energy for complexes of aromatic hydrocarbons with p-xylene us. the electron affinity of the aroma tic hydro carbon.
sumptions is that correlations between AFO and IB and E.& do exist. At best, probably only a rough correlation could be expected and it is remarkable that the correlation is as good as it is. Figure 5 demonstrates the correlation between AF" and the ionization potential for the series of methyl-
I
.22
(e.v.Y'
Fig. 7.-Change in st'andardfree energy for aromatic hydrocarbon complexes with polymethylbenzenes os. (IB - E A - 3)-'.
benzenes as donors with anthracene. Figure 6 illustrates the relat'ionship between AFO and the electron affinities of the series of acceptors with p-xylene. The horizontal lines above and below the points indicate the standard deviation of AFO. This mas obtained from the propagation of error in K upon AFO. The error in K was obtained from the least squares solution as discussed by Deming. Actually, there should be some curvature in both of the correlat#ions,but since all three constants could not be determined and since there is some question about their constaiicy, straight lines have been drawn to show that the general trend is followed in agreement' with eq. 19. It should be noted that AFo increases negatively as the I B is decreased in agreement with eq. 19. In ail analogous manner AFO increases negatively as the EA is increased, again in agreement with eq. 19. These two correlations, in agreement with eq. 19, appear to indicate rather conclusively that charge-transfer stabilization a t least plays some role in complex formation. Figure 5 is comparable to the correlation given between the stabilities for the series of complexes formed by the methylbenzenes with tetracyanoethylene and the ionization potential of the donors given by Merrifield and Phill i p ~ . If ~ their ~ plot is extended by using recent data on the ionization potentials of the higher polymethylbenzenes, definite curvature can be established in agreement with eq. 19. Considering the range of IB - Ea values for these complexes, one cannot accurately establish all three const'ants C1, C2,C3. However, one can make a reasonable guess a t Czsince it primarily involves the coulombic attraction of the ions in the dat'ive state. Hast'ings, (23) It. ilO5S).
E:. hlerrifield and TT. D. Phillipe, .I. A m . Chrm.
S o r , . , 8 0 , 2778
Oct., 1963
TNTERACTIOX
BETWEEN
METHYLBENZENES AND POLYCYCLIC HYDROCARBONS
ul., have estimated this constant as 4 e.v. for a series of complexes with 12. Since the polycyclic aromatic hydrocarbons are much larger than 1 2 , we have reduced this estimate to 3 e.v. Actually, the selection of this constant is not critical within two electron volts and does not seriously affect the following correlation. Figure 7 shows a plot of AFO vs. l/(IB- EA-. C2) with Cz equal to 3 e.v. From the theory, this should be a straight line if the constants for the series of acceptors with p-xylene are the same as the constants for the series of donors with anthracene. The correlation is remarkably good considering the approximations involved in the constant terms. As stated previously, one could examine this relationship more critically if the associated charge-transfer frequencies were available. Also, it shlould be stated that the results of this work do not directly contradict the work of authors in other fields. I n a paper by Porter and WilkinsonZ4it was stated that there was no evidence of complex formation between phenanthrene and naphthalene as examined by absorption spectra and hence these authors eliminated this process as a mechanism of energy transfer. Taking into account the concentrations a t which they were working and assuming that naphthalene would fall into the same series as the methylbenzenes, it is estimated that only 03% of the phenanthrene would exist as a complex. This small magnitude could hardly be detected from spectrophotometric measurements. Finally, this work supports the conclusion of Robertsen, Babb, and h/8atsen2jconcerning the absence of a self complex of benzene since benzene has both a high ionization potentia,l and a low electron affinity and would have a very fimall association constant. et
Conclusions 1. The absorption data as a function of the donor concentration satisfy a 1: 1complex formation equilibria and give a nonzero equilibrium constant for most comlbinations studied. Of course this does not uniquely de(24) G. Porter and F. Wilkinson, ”Luminescence of Organic & Inorganio :Naterials,” John Wiley and Sons, Inc., New York, N. Y., 1962, pp, 132-142. (25) W. W. Robertsen, S. E. Babb, a n d F. A. Matsen, J. Chem. Phys., 26, 3 6 7 (1957).
2207
fine the equilibria since more complicated equations could fit the data equally well. 2. The correlation of the AFO with the ionization potential of the donor and the electron affinity of the acceptor is evidence that the complex formation may be due to the charge-transfer stabilization as described by Mulliken. The correlation also further justifies the previous experimental measurements of the electron affinities of the aromatic hydrocarbons. 3. Based upon the previous conclusions, it can be stated that the aromatic hydrocarbons can act as acceptors in charge-transfer complexes. Implications and Future Work Masonz6 has suggested that the carcenogenicity of the aromatic hydrocarbons arises from their electron acceptor ability and this suggestion has been strengthened by Lovelock, et ul.27~28This work supports the suggestion and indicates that a possible mechanism for carcenogenicity may be some complex formation in which the hydrocarbon acts as an acceptor. Evidence for this was recently presented by Allison, et ~ 1 . ~ At present, work is in progress to obtain experimental electron affinities of stronger acceptors and to correlate these values with AHo and h v C p With both of these experimental parameters and the known ionization potentials, the coiistants in eq. 19 can be examined more critically. Acknowledgment.-This work mas supported by grants from the Robert A. Welch Foundation (held by W. E. Wentworth) and by the Tobacco Industry Research Committee (held by R. S. Becker). The authors are most grateful for this support. We also wish to express our appreciation to Dr. R. S. Becker for the polynuclear aromatic hydrocarbon samples and for his general interest, support, and criticism of the work. Finally, we thank Mr. R. G. Allen for his most cooperative effort in programming the least squares problem on the IBM 7090. (26) (27) (28) (29)
R . Mason, Discussions Faraday Soc., 27, 129 (1959). J. E. Lovelock, Nature,189, 729 (1961). J. E. Lovelock, R. S. Becker, and A. Zlatkis, %bid., 193, 540 (1962). A. C. Allison, M. E. Peover, and T. A. Gough, ibid.,197, 764 (1968)
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