J . Phys. Chem. 1990, 94, 2187-2190
0
2
4
6
8
1
0
106[AQSI I M
Figure 11. Plots of loll vs [AQS] for the fluorescence quenching of TMPyP (0)and TAPP ( 0 )by AQS at room temperature. The porphyrin concentrations were 5 X 1od M. The fluorescence intensities were followed at 704 nm.
to know the abilities of TMPyP to form self-aggregate(s) and molecular complexes in aqueous media. In spite of various attempts to clarify the self-association phenomenon of TMPyP in water,’-7 the direct evidence for dimerization of this porphyrin has not been reported as yet. The present study provides several experimental aspects that support the dimer model of TMPyP. ‘H N M R clearly indicates the formation of the self-aggregates of TAPP having fast exchange rates in concentrated aqueous solution. Since a Lambert-Beer’s law can be applied for the Soret band of TAPP in the concentration range 2 X to 5 X M, the monomer should be the main species in dilute aqueous solution. Partial neutralization of the positive charges localized on the trimethylammonium groups by added NaCl promotes the asso(25) Pasternack, R. F.; Gibbs, E. J.; Villafranca, J. J. Biochemistry 1983, 22, 2406. (26) Marzilli, L. G.; Banville, D. L.; Zon, G.;Wilson, W. D. J. Am. Chem. SOC.1986, 108, 4188. (27) McKinnie, R. E.; Choi, J. D.; Bell, J. W.; Gibbs, E. J.; Pasternack, R. F. J . Inorg. Biochem. 1988, 32, 207.
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cation of the TAPP molecules. This behavior of TAPP is similar to that for the anionic water-soluble p o r p h y r i n ~ , ~ , whose ~J~J~ negative charges are localized on the substituents at the meso positions. The formation of the 1:l complex of TAPP and PFl also supports the weak ability of TAPP to associate spontaneously. Meanwhile, most experimental results obtained for TMPyP and TOPyP can be understood only by the dimer model for these cationic porphyrins. The stoichiometries determined for the TMPyP- and TOPyP-PFI complexes (2:l) as well as that for the TMPyP-BSA complex (4: 1) strongly suggest the extremely stable dimers of these pyridinium porphyrins. Since the pyridinium substituent rotates around the C-C single bond at the meso position, the positive charge on the substituent can pour onto the porphyrin ring when the pyridinium ring becomes coplanar to the porphyrin ring. The dipole thus generated seems to induce the dipole of another porphyrin having a well-extended conjugation leading to the strong T-T interaction and reduction in the electrostatic repulsion between these two porphyrin rings. The fact that the association constants for complexation of TMPyP and TOPyP with an cationic dye, PF1, are much larger than that of TAPP with PFl supports this assumption. Appearance of the well-separated fluorescence Q bands of TOPyP upon addition of NaCl and/or AQS is also explained by the dimer model for this porphyrin. If the coalescence of the fluorescence Q bands of TMPyP and TOPyP is ascribed to some electronic effect, the separation of the Q bands should also be observed for TMPyP when NaCl and/or AQS is added, because the electronic structure of the porphyrin ring of TMPyP is the same as that of TOPyP. The novel behavior of TOPyP in water containing NaCl and/or AQS can be interpreted in terms of the dissociation of the TOPyP dimer to the monomer due to the effects of the hydrophobic octyl groups. Although we could.not derive the direct evidence for the TMPyP dimer model from the present study, all results suggest strongly that TMPyP in water exists in its dimer form even at very low concentrations.
Acknowledgment. The present study was supported by a Grant-in-Aid for Science Research No. 62470082 from the Ministry of Education, Science and Culture. We are indebted to H. Abe and H. Takagi for their experimental assistance. Registry NO.TOPyP, 124891-59-4;TOPyP(I), 104207-96-7;TMPyP, 92739-63-4; TAPP, 92739-64-5; PFI, 92-62-6; AQS, 13 1-08-8; ZnTMPyP, 28850-44-4.
Evaluation of Group Electronegativity by Pauling’s Thermochemical Method Dipankar Datta* and S.Nabakishwar Singh Department of Chemistry, Manipur University, Imphal 795 003, India (Received: March 13, 1989) Pauling’s bond energy equation, D(H-G) = [D(H-H) D(G-G)]’12 + K(Ax)*(l), is applied to some 28 HG molecules, where H is hydrogen and G any group of atoms, to calculate the electronegativity (xc) of a group. In eq 1 D(H-G), D(H-H), and D(G-G) are the energies of the H-G, H-H, and G-G bonds, respectively, Ax = IxH- xGI,and K is a constant. The value of K is optimized to 25 in order to obtain maximum amount of correlation ( r = 0.960) of the xc values for some 19 groups/atoms with the ‘Jcc(ortho-ipso) coupling constants in monosubstituted benzenes. These ‘Jcc constants form a very good experimental scale for group electronegativity. With Inamoto’s ,in scale, derived experimentally, the xGvalues for some 20 groups/atoms correlate quite satisfactorily (r = 0.936). However, Ax for two groups, CH(OH)CH3 and Si(CH,),, out of the 28 groups considered are found to be imaginary, for which the bond energy data are held partially responsible. The effects of various bond-determining factors on the xGvalues evaluated by eq 1 are examined critically. It is concluded that, when the order of the G-G bond is 1 as in H-H or H-G, the method yields a reasonable value for the electronegativity of a group.
The electronegativity of an atom was first defined by Pauling.] Subsequently, extension of this concept to a group of atoms gave
rise to a term “group electronegativity” and various approaches to evaluate this parameter with various degrees of success (for
0022-3654/90/2094-2187$02.50/00 1990 American Chemical Society
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Datta and Singh
TABLE I: Bond Dissociation Energy Data, Group Electronegativity, and Various Measures of Group Electronegativity“ Yingst’s Mullay’s Datta’s group G D(H-G) D(G-G) IAxI XG datab datac datad 1. H 103.258 103.258 2.1h 2. CH, 105.1 90.4 0.58 2.68 2.64 2.32 2.80 98.2 82.2 0.49 2.59 2.71 2.35 2.77 3. C2H5 0.44 95.1 79.0 2.54 2.60 2.38 2.75 4. i-C3H7 93.2 71.2 0.55 2.65 2.68 2.41 2.76 5. t-CdH9 4.17 119.0 51.0 1.36 3.46 3.53 3.97 6. OH 104.4 37.6 3.40 4.03 4.19 1.30 7. OCH, 37.9 1.29 3.39 104.2 8. OC2HS 103.4 37.1 1.29 3.39 9. O-n-C,H7 104.7 37.1 1.31 3.41 IO. O-i-C,H7 3.43 105.5 36.4 1.33 1 1 . O-sec-C4H9 1.31 105.5 38.0 3.41 12. O-t-CdH, 3.38 I .28 102.3 36.4 13.O-CH,C(CH,), 4.18 105.8 30.4 3.51 1.41 4.33 14.02CCH3 1.42 106.4 30.4 3.52 15. 02CC2Hy 30.4 3.51 1.41 105.9 16. 02C-n-C3H7 2.44 2.89 2.70 87 68.4 0.34 17. C H O 67.4 2.42 0.32 2.72 2.93 86.0 18. COCH, 0.40 66.4 2.50 86.9 19. COC,Hs 94 2.45 0.35 2.75 2.50 80.2 20. CHiOH 93 84.9 21. CH(OH)CH, i 2.42 66 0.55 2.65 90.1 22. SH 0.44 2.54 72 23. SR 91 3.58 3.10 3.17 65.8 24. NH2 107.4 1 .oo 3.15 13.6 3.93 3.38 1.28 4.08 25. NO2 78.3 4.73 4.0h 4.73 26. F 3.0h 27. CI 2.85 2.58 3.10 3.10 98.7 0.48 28. CF, 106.7 3.18 21 1.08 75.7 29. NF2 1.76 1 .80 1.97 74 0.34 90.3 30. SiH, 80.5 90.3 31. Si(CH,), j 1 .Oh 1.01 1.68 32. Li
iJ
‘JC2
0.00 1.07 1.09 1.5 1.9 9.7 1 1.09
2.00 2.14 2.15 2.15 2.16 2.79 2.82
12.78
2.80
2.0 1.9
2.61 2.39
1.65
2.22
4.2 3.9k 5.2 11.43 14.84 9.21 3.57
2.17 2.16 2.47 2.75 3.1 2.37 2.47
-6.5
1.79‘ 1.79 1.15
‘The meanings of the symbols are same as in the text. The bond energy data are given in kcal/mol and taken from ref 8 unless otherwise specified. The x values are given in Pauling’s unit. bFrom ref 7. CFrom ref 4. dFrom ref 5. eIn Hz; from ref 10. fFrom ref 1 1 . BFrom p A32 of ref 21. From ref 1 . jlmaginary. For SCH3. ‘For Si(CH,),.
brief reviews on this aspect, see ref 2-5). Originally, Pauling adopted an thermochemical method for measuring atomic electronegativity. It is surprising to note that, although there have been a number of methods to evaluate group electronegativity, Pauling’s own method has never been tested seriously. While the lack of appropriate thermochemical data may be one of the reasons, a more serious one is probably the fact that the connection between Pauling’s original definition of electronegativity and his thermochemical scale has never been clearly understood. In this context it should be mentioned that Wells2 has tried to justify Pauling’s approach in terms of qualitative MO theory. However, in the past only Constantinides6 and Yingst’ attempted to use Pauling’s method with slight variations for calculating the electronegativity of a very few groups. But apparently no appreciable attention has ever been paid to their results. Herein we extend Pauling’s original method to calculate electronegativity of some 26 groups and examine the workability of the method critically. According to Pauling,’ D(P-Q), the energy of the P-Q bond in a diatomic molecule PQ is related to the bond energies D(P-P) and D(Q-Q) of the molecules PP and QQ, respectively, by eq 1, D(P-Q) = [D(P-P)D(Q-Q)]’/* + K(Ax)* (1) where K is a constant and Ax the difference between xpand xQ, the electronegativities of the neutral atoms P and Q. The value of K prescribed by Pauling for diatomic molecules is 30. We assume that a similar relation where the value of K may be different than 30 holds good for the polyatomic molecules also (1) Pauling, L. J. Am. Chem. SOC.1932, 54, 3570. The Nature of the Chemical Bond, 3rd ed.; Oxford and IBH Publishing Co.: New Delhi, 1967; Chapter 3. (2) Wells, P. R. Prog. Phys. Urg. Chem. 1968, 6, 11 1. (3) Bratsch, S. G . J. Chem. Educ. 1985, 62, 101. (4) Mullay, J. J . Am. Chem. SOC.1985, 107, 7271. (5) Datta, D. Proc. Indian Acad. Sei. (Chem. Sei.) 1988, 100, 549 (6) Constantinides, E. Proc. Chem. SOC.1964, 290. (7) Yingst, A. Chem. Commun. 1965, 480.
I j5
25
35
K
Figure 1. Variation of the coefficient ( 7 ) of correlation between the X G values and the ‘JCccoupling constants with the constant K i n eq 1 (see text). The curve drawn is not least-squares fitted.
at least in the cases where one of the entities P and Q is an atom and only a u bond exists between P and Q. For this purpose, the HG molecules where H is hydrogen and G any group of atoms seem to be best suited since H does not have any pseudo-7r orbitals which can lead to hyperconjugation with the group G. For a number of H G molecules, we have collected the data (Table I) on D(H-G) and D(G-G) from a recent compilation of bond dissociation (homolytic) energies by Golden and McMillen.* _
_
~
_
~
_
~
~
(8) McMillen, D F ; Golden, D M Annu Rev Phys Chem. 1982, 33, 493
The Journal of Physical Chemistry, Vol. 94, No. 5 , 1990 2189
Group Electronegativity by Pauling’s Method
1s.o
-
i0.0
-
3s
i 4
Jic
5.0.
0.0
2s
-
-so -
1.0
1.0
2.0
3.0
4D
% 10
90
%
Figure 3. Correlation between the xG values and Inamoto’s i scale (see Table I for data).
40
3.0
Figure 2. Final correlation between the xo values and the ‘Jcc coupling constants (see text).
Using eq 1 with a certain value of K ( K was varied around 30), At present we have calculated xGfor some 26 groups (Table the IJcc(ortho-ipso) coupling constants in the monosubstituted benzenes (Table I) are regarded as the best experimental measures of group e l e c t r o n e g a t i ~ i t y . ~Hence ~ ~ J ~ these coupling constants were chosen as the reference data set for optimizing K . As seen from the Figure 1, K = 25 gives the best correlation ( r = 0.960; r;! = 0.922) of the XG values with the reference scale for some 19 groups/atoms (xF and xcl were kept untouched while varying K ) . Thus the final equation used for calculation is I).99’0
D(H-G) = [D(H-H) D(G-G)]’12
+ 25(Ax)’
(2)
The xGvalues given in the Table I correspond to eq 2. In a way the very good correlation obtained for the various xGvalues with the ‘Jccconstants (Figure 2) justifies our assumption regarding the applicability of eq 1 in the HG molecules. So far, of all the scales available for group electronegativity, only the scales of Mullay4 and DattaS correlate well with these coupling constants. And the amounts of correlation achieved by them (r = 0.9404and r = 0.9695)are comparable to the present case. However, it should be mentioned that for two groups CH(OH)CH3 and Si(CHJ3 the method encounters imaginary values of Ax (see Table I). It is very difficult to say whether the method fails or the bond energy data are in error. Even an increase in the D(H-G) values for the groups in question by 1 kcal/mol would yield real values of Ax. Similar problems were faced by Pauling himself while calculating atomic electronegativities, but for this he held the bond energy data as partially responsible.’ (9) Electronegativity of a species cannot be calculated directly from eq
1-it only gives [Ax].To find out the sign of Ax with respect to hydrogen, we have considered the results of the recent ab initio calculations by Marriott et a1.I0of group charges in the HG molecules at 6-31G1 level. For example, since in H-CH,the methyl group bears negative charge, it is more electronegative than H and thus xCH = xH + IAxI. Where ab initio data are not available we have been guided by the general notion of electronegativity. (10) Marriott, S.; Reynolds, W.F.; Taft, R. W.; Topsom, R.D. J . Org. Chem. 1984, 49, 959.
Another experimental scale, claimed to represent group electronegativity, is Inamoto’s “i” scale.” The i scale is again derived from extensive correlations of the N M R data for a wide variety of chemical groups. With this i scale (Table I), for some 20 groups/atoms the XG values correlate quite satisfactorily ( r = 0.936; Figure 3). Almost same amount of correlation (r = 0.935) is produced by Datta’s scale,s but Mullay’s scale correlate much better ( r = 0.975)4 with the i values. The basic assumption in applying eq 1 to the HG molecules is that while the nature of the bonds H-H and G-G is purely covalent, that of H-G is largely ionic; the covalent part in H - G with respect to H-H and G-G is likely to be compensated in the difference between D(H-G) and the geometric mean of D(H-H) and D(G-G); Le., A = D(H-G) - [D(H-H) D(G-G)]’12 and the uncompensated bond energy difference is correlated with the electronegativity difference Ax assuming ionic interactions to be the only origin of it. But electronegativity difference is only one of the bond-determining factors. To mention a few, among the other factors are matching of orbital overlaps and “local hardnesses”, steric repulsions, etc. For very good discussions on the various factors responsible for a chemical bond, see ref 12-14. Here it is necessary to make some comments upon the effect of the matching of local hardnesses on bonding. Hardness of a chemical species is defined as (I - A)/2, where I and A are respectively the ionization potential and electron affinity of the neutral specie^.^^^^^ For closed-shell species this quantity can be correlated with the energy gap between the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO).” Unlike electronegativity [=(I + A)/2], hardness is not the same at every point in the species; Le., hardness allows for local values-“local hardness”.”J* From the recent (11) Inamoto, N.; Masuda, S . Tetrahedron Lett. 1977, 3287. Inamoto, N.; Masuda, S.; Tori, K.; Yoshimura, Y . fbid. 1978, 4547. (12) Fliszar, S. J . Am. Chem. SOC.1980, 102, 6946. (13) Kutzelnigg, W. Angew. Chem., Znr. Ed. Engl. 1984, 23, 272. (14) Pearson, R. G. J . Am. Chem. Sac. 1988, 110, 7684. (15) Parr, R. G.; Pearson, R. G . J . Am. Chem. SOC.1983, 105, 7512. (16) Datta, D. J . Phys. Chem. 1986, 90, 4216. (17) Pearson, R. G. Proc. Natl. Acad. Sci. U.S.A. 1986, 83, 8440. (18) Berkowitz, M.; Ghosh, S. K.; Parr, R. G. J . Am. Chem. Sor. 1985, 107, 6811. Berkowitz, M.; Parr, R. G. J . Chem. Phys. 1988, 88, 2554.
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elucidations of the principle of hard and soft acids and bases (HSAB) by P e a r ~ o n , ’ ~it*follows ’~ that correct matching of local (at the reaction sites) hardness between two interacting chemical species (an acid and a base), X and Y, gives rise to an extra stability of the X-Y bond. However, this factor becomes particularly important for heterolytic dissociation of X-Y. Since we are dealing with homolytic dissociations here, this factor is not of much consequence to our present study. We understand that the effect of any factor not compensated in the difference A will affect the value of Ax. There is a postdissociation phenomenon which affects the homolytic bond dissociation energy considerably-the resonance stabilization of the radicals generated in the process. Such stabilization can be quite substantial in radicals like benzyl, carboxyl, or nitro, where delocalization of the unpaired electron over several atoms is possible through conjugation, bringing about equivalent amount of lowering in the D However, this factor is roughly cancelled in the difference A. While probably most factors will be nullified to a reasonable extent as we consider the difference A, it will be appropriate to point out only those factors (others than electronegativity) likely to contribute to the value of A x . First, the bond order (bo) of the G-G bond may not be one ,~ as of the H-G bond. For example, as pointed out by Y i n g ~ tthe c-c bond in NC-CN or H5C6-C6H5has considerable amount of double bond character. The increase in the bo of the G-G bond would result in a xG value much lower than the actual one. The groups potentials in this regard have not been included in our study. Second is the steric effect, a result of nonbonded interactions. This effect which depends on the bulk of the groups is not significant for a H - G bond but may be quite so for a G-G bond as in (CH3)3C-C(CH3)3. Steric repulsions generally destabilize a G - G bond and thereby would increase the value of xG. However, steric factors are known to change the hybridization of nitrogen in NR,R2R3(R = alkyl group) from sp3 to sp2 making the N-R bond stronger than expected (back strain).2’ Hence caution should be used in appraising this effect. Lastly there is another effect recognized recently to influence the bond energy. In inorganic chemistry it is known as symbiotic effect and in organic chemistry it is anomeric, clustering, or geminal effect.I4 In Pearson’s language this “refers to the stabilization caused by adding several hard substituents to the same” (19) Pearson, R. G.Inorg. Chem. 1988, 27, 734. (20) Halpern, J. Acc. Chem. Res. 1982, 15, 238. (21) Huheey. J. E Inorganic Chemisrry, 3rd ed.; Harper and Row: New York, 1983: pp 298-299
Datta and Singh central atom. Two explanatory examples from organic chemistry are given below. 4CH,F(g) = 3CH4(g) + CF4(g)
AHo = -63 kcal/mol
4CH3CH3(g) = 3CH4(g) + C(CH,),(g) AH = -13 kcal/mol The effect is not clearly understood at present but definitely not accounted for in eq 1. As revealed by the goodness of correlation in the xc values with the ].Iccconstants, in the absence of the bo problem (vide supra) the contribution of the various factors uncompensated in the difference A toward Ax has not been significant. In fact certain expected trends are preserved. For example, though the ’.Icc values (Table I) show the order of the electronegativities of the alkyl groups as CH, < C2H5 < i-C3H7< t-C4H9, the opposite is always expected on the basis of the knowledge of inductive effects displayed by these groups; from our present study the order is found to be CH, > t-C4H9> C2H5> i-C3H7where only t-C4H9 is an exception. Though eq 2 correctly gives xOH> xOR,the inductive effect of R seems to be levelled off for all R groups. Similar levelling off is also observed for the three 0 2 C R groups (R = CH,, C2H5,and n-C,H,) studied. The trend observed for COH, COCH3, and COC,H5 is remarkably correct. Since CH3 is an electron-donating group and C6H5an electron-withdrawing > xCPH> xCOCH3. While the broad one, one truly expects features of group electronegativity are reflected, in general, correctly, the only exception in our study seems to be the CF, group (xG = 2.58) which is found to be less electronegative than even the methyl group (xG = 2.68). This is not really acceptable, but we note that, in Datta’s scale,5 which is based on ab initio calculations, electronegativity of CF3 is comparable to that of CH3 (see Table 1). It is worthwhile mentioning that Pauling’s bond energy equation has been criticized earlier by several workers in some other context^.^^^^^^^^ However, from our present study we conclude that Pauling’s thermochemical method can be applied to the chemical groups as well with minor modification (Le., changing K in eq 1) for evaluating group electronegativity at least with respect to hydrogen provided the bo problem mentioned in an earlier section is carefully avoided.
Note Added in Proof. As discussed in an earlier section, steric strain in t-C4H9-t-C4H9 may be responsible for increasing the x , . value ~ ~and~ hence ~ for the deviation of the t-C4H9group from the expected trend (vide supra). (22) Pearson, R. G. Chem. Commun. 1968, 65. (23) Evans, R. G.; Huheey, J. E. J . Inorg. Nucl. Chem. 1970, 32, 373.
