6902
J. Phys. Chem. 1988, 92,6902-6907
EIechton Donor-Acceptor Propertks of Porphyrins, Phthalocyanines, and Related Rlng Chetates: A lblecular Orbital Approach C. Fierro, Alfred B. Anderson, and D. A. Scherson* Case Center for Electrochemical Sciences and The Department of Chemistry, Case Western Reserve University, Cleveland, Ohio 44106 (Received: March 4, 1988; In Final Form: June 2, 1988)
Tetraazaporphyrin is used as a primitive fragment to examine theoretically the electron donoracceptor properties of a variety of highly conjugated planar ligands of the porphyrin and phthalocyanine type. Calculations employing the atom superposition electron delocalization MO theory indicate that for iron macrocycles the energy gap between the occupied d,, dYzorbitals and the empty e,(**) ring orbitals plays a key role in controlling the extent of metal to ring back-bonding. This provides a framework for a rational search of structural modifications of chelating ligands that may afford an optimized activation of axially coordinated A electron acceptor molecules such as dioxygen, nitric oxide, and carbon monoxide. The higher reactivity of iron(I1) porphyrins toward such A acceptor species, compared to iron(I1) phthalocyanine, is attributed to the longer Me-N bond length in the porphyrins, which leads to less metal to ligand back-bonding and thus to more electron charge at the metal center.
Introduction A number of reports have appeared in the literature over the past two decades describing the ability of certain transition-metal phthalocyanines and porphyrins to catalyze the electrochemical reduction of di~xygen.'-~ On the basis of results obtained by several groups, the most active macrocycles for O2 reduction appear to be those involving iron or cobalt centers. Among these, iron tetrasulfonated hthalocyanine&* (FeTsPc), iron tetrapyridinoporphyrazine,Jand certain face-to-face metal porphyrins synthesized independently by Collman et a1.5and Chang et ala6 are especially interesting because of their ability to reduce dioxygen directly to water without any apparent generation of peroxide. This exceptional behavior is unlike that of most other Fe and Co macrocycles which either produce significant amounts of peroxide as an intermediate or are unable to promote the reaction beyond the initial two-electron transfer, yielding peroxide as the final product. Although arguments have been given' to explain the relative reactivity of transition-metal elements toward 02,NO, and other first-row diatomic molecules in terms of the energy difference between the d orbitals and the diatomic A* orbitals, a rationalization of the way in which the chelate ring structure may influence the ability of complexes to activate such species has not been pursued. This seems of importance since highly conjugated planar chelates have electron accepting orbitals which can compete with electron acceptor molecules such as axially coordinated 02,CO, and N O for electron density from the metal. This paper will apply the principles of molecular orbital theory to analyze the electronic properties of iron macrocycles as calculated by using the atom superposition electron delocalization (ASED-MO) theory.8 The use of a simple semiempirical ap(1) For a review see Tarasevich, M.; Sadkowski, A,; Yeager, E. In Comprehensive Treatise of Electrochemistry; Conway, B., et al., Eds.; Plenum: New York 1984; Vol. 7. (2) Murray, R. W. Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1984; Vol. 13 and references therein. (3) Yeager, E. Electrochemistry in Industry: New Directiom; Landau, U., Yeager, E., Kortan, D.,Eds.; Plenum: New York, 1982. (4) (a) Zagal, J.; Bindra, P.; Yeager, E. J . Electrochem. SOC.1980,127, 1506. (b) Sen, R.; Zagal, J.; Yeager, E. Inorg. Chem. 1977, 16, 3379. (c) Zagal, J.; Sen, R. K.; Yeager, E J. Electroanal. Chem. Interfacial Electrochem. 1977,83, 207. ( 5 ) (a) Durand, R. R.; Bencosme, C. S.; Collrnan, J. P.; Anson, F. C. J. Am. Chem. SOC.1983, 105, 2710. (b) Collman, J. P.; Bencosme, C. S.; Durand, R. R.; Kreh, R. P.; Anson, F. C. Ibid. 1983,105,2699. (c) Collman, J. P.; Anson, F. C.; Barnes, C. E.; Bencosme, C. S.; Gliger, T.; Evitt, E. R.; Kreh, R. P.; Meier, K.;Pettman, R. B. Ibid. 1983, 105,2694. (d) Durand, R. R.; Collrnan, J. P.; Anson, F. C. J. Electroanal. Chem. Interfacial Electrochem. 1983, 151, 289. (6) Liu, H. Y.; Weaver, M.;Wang, C.; Chang, C. J. Electroanal. Chem. Interfacial Electrochem. 1983, 145, 439. (7) Hoffmann, R.; Chen, M. M.-L.; Thorn, D. L. Inorg. Chem. 1977, 16, 503.
0022-365418812092-6902$01 SO10
proach is a highly desirable first step in rationalizing the bonding and the overall chemical characteristics of these systems. An excellent illustration of the power of a closely related technique may be Seen in the work of Hoffmann and co-workers, who have made extensive use of orbital overlap and symmetry arguments and perturbation theory based on extended Hiickel calculations9 to gain insight into molecular structure and reactivity. Particular emphasis will be placed on the analysis of the ring-iron center orbital interactions and on the modifications in the electronic charge density at the metal center which are introduced by peripheral, axial, and ring substituents on a primitive macrocycle fragment. Although attention has been focused only on iron macrocycles, the methodology employed can be easily extended to encompass other metal and nonmetal centers. The atomic parameters involved in the calculations are those reported in earlier studies (see Table I). The atomic coordinates of the compounds examined, obtained from X-ray data, have been slightly modified so as to achieve a D,,,, symmetry, and reasonable bond angles and distances are chosen when necessary without attempting any structure optimization. Model Systems for Metal Porphyrins and Phthalocyanines Due to the fairly large size of macrocyclic molecules, several authors have introduced simplified model compounds in order to investigate theoretically the physicochemical properties of iron porphyrin ( F e P la) iron phthalocyanine (FePc; l b ) and their derivatives both in monomeric and polymeric forms.
la
lb
Sabelli and Melendres,14 for instance, considered the 16 ?r electron system 2 (obtained upon replacing the benzene fragment of each isoindole moiety in Pc, and the ethylene fragment in P, by two hydrogen atoms), to examine the corresponding substituted compounds. On the basis of CNDO calculations, it was concluded that the bridge atoms ( N in Pc and C in P) are mostly responsible (8) (a) Anderson, A. B. J. Chem. Phys. 1974,60,2477. (b) Anderson, A. B. J. Chem. Phys. 1975, 62, 1187. (9) Hoffmann, R. Acc. Chem. Res. 1971, 4, 1. (10) Rhodin, T. N.; Brucker, C. F.; Anderson, A. B. J. Phys. Chem. 1978, 82, 894. (11) Clementi, E.; Raimondi, D. L. J. Chem. Phys. 1963, 38, 2686. (12) Lotz, W. J. Am. Opt. SOC.1970, 60, 206. (13) Anderson, A. B.; Kang, D. B. Inorg. Chem. 1984, 23, 1170. (14) Sabelli, N. H.; Melendres, C. A. J. Phys. Chem. 1982, 86, 4342.
0 1988 American Chemical Society
The Journal of Physical Chemistry, Vol. 92, No. 24, 1988 6903
Electron Donor-Acceptor Properties of Macrocycles
TABLE I: Parameters Used in the Calculations: hincipal Quantum Number ( n ) , Ionization Potential (eV), Orbital Exponents 6 (au), and Reswctive Coefl'icients (d Onlv)" ~~
P
S
d
atom
n
IP
5
n
IP
5
n
IP
t
Fe F 0
5.83 11.42 13.62 12.44 11.26
1.4 2.55 2.2266 1.917 1.618
10.5
1.8
1
1.7 2.5639 2.146 1.9237 1.658 1.2
4 2 2 2 2
H
9.37 31.85 28.48 20.33 20.00 13.6
3
C
4 2 2 2 2
N
5.35
0.6678
0.5366
"The exponents for iron and carbon are from Rhodin, Bruker, and Andersonto and those for the other atoms are from Clementi and Raimondi." The value of 1.2 for H is a standard value in extended Hiickel calculations. Ionization potentials are based on the data of Lotz.I2 The ionization potentials in the case of iron are increased by 1.5 eV to approximate the results of charge ~elf-consistency.~~
for the differences in the charge distribution between porphyrins and phthalocyanines. In the porphyrin case, more than 82%of the charge was localized in the pyrrole ring, leading the authors to suggest that reasonable results could be obtained by considering the model compound shown in 3. This simplified structure was
L
used by Goddard et al.15 in an a b initio study of the preferred orientation of O2in an axially coordinated porphyrin adduct and by Hoffmann and co-workers in an extended Hiickel study of metalloporphyrin dimers.16 The latter authors employed the complete 18 7 electron model whenever the metal-7-ring interaction was believed to play an important role. The conduction band properties of one-dimensional metal phthalocyanine polymers were examined by Whangbo and Stewart" and Canadell and Alvarez18using 2 as the basic model in a tight-binding band structure analysis. The stabilization of the highest occupied molecular orbital (HOMO) of the ring, when 2 interacted with the 7 and T* MO's of either a benzene or ethylene fragment to yield a metal phthalocyanine l b or tetraazaporphyrin 4, was used as a criterion for estimating the error
and manganese porphyrin4ioxygen adducts. This model has not been widely used because of difficulties introduced in the orbital analysis by the loss of Dqhsymmetry, which lifts the degeneracy of the d,, dyz orbitals. Fragment Orbital Analysis Tetraazaporphyrin (TAP) 4 was selected in this study as a primitive model for the analysis of a variety of macrocycles. This same structure was used recently by Anderson et a1.21in a M O analysis of the cyclic voltammetry features of stacked-ring silicon phthalocyanines.22 The atomic coordinates for 4 have been taken from the X-ray study of F ~ P by c ~omitting ~ the appropriate moieties and averaging the bond distances so as to achieve a D4h symmetry. General Symmetry and Overlap Considerations A symmetry analysis of the metal-ring orbital interactions has been presented earlier in the literature.16 As sketched schematically in 6 (only the contribution localized on the closest
a 4
in the band structure associated with this simplified model. Using the discrete variational local exchange (DV-Xa) technique, Ratner, Marks, and co-w~rkers'~ found similar results by using the complete phthalocyanine molecule. A different model structure, shown schematically in 5 was
H-
A *
H
2-
-
--H
'
y
k 5
introduced by Newton and HallZoto investigate a series of iron (15) Goddard, B. D.; Goddard, 111, W. Proc. Natl. Acad. Sci. U S A . 1977, 74, 1315. (16) Tatsumi, K.; Hoffrnann, R. J . Am. Chem. Soc. 1981, 103, 3328. (17) Whango, M.-H.; Stewart, K. R. Isr. J . Chem. 1983, 23, 133. (18) Canadell, E.;Alvarez, S. Inorg. Chem. 1984, 23, 573. (19) (a) Ciliberto, E.; Doris, K. A.; Pietro, W. J.; Reisner, G. M.; Ellis, D. E.;Fragala, J.; Herbstein, F. H.; Ratner, M. A.; Marks, T. J. J. Am. Chem. Sw.1984, 106,7748. (b) Pietro, W. J.; Marks, T. J.; Ratner. M. A. J. Am. Chem. SOC.1985, 107, 5387.
nitrogen is depicted), the ring bl,, bp, e,(*) (or e,(**)), and alg orbitals have the appropriate symmetnes to mix with the respective
metal dA9, d,, d(*) (or d,, ,,=), and d s orbitals. Based on the directional character of the d+? and the bl, orbitals, the u bonds ~~~
~
(20) (a) Newton, J. E.; Hall, M. B. Inorg. Chem. 1984, 23, 4627. (b) Newton, J. E.; Hall, M. B. Inorg. Chem. 1985, 24, 2573. (21) Anderson, A. B.;Gordon, T. L.; Kenney, M.E. J. Am. Chem. Soc. 1985, 107, 192. (22) (a) Mezza, T. M.; Armstrong, N. R.; Ritter, G. W.; Iafalacie, J. P.; Kenney, M. E. J. Electroanal. Chem. Interfacial Electrochem. 1982,137,227. (23) Kirner, J. F.; Dow, W.; Scheidt, R. W. Inorg. Chem. 1976, IS, 1685.
6904 The Journal of Physical Chemistry, Vol. 92, No. 24, I988
Fierro et al.
J
-5
-
-6
-
-7
-
?)I'
> P) Y
*
-6
!-
-7
-
-0
-
Y
\
-9 -
9
(3 - 1 0 -
a
w w
-8
-
:-"" I ,
Q)
Y
-11-12
-
-1 3
-
-14
-
-15
-
> (3
CT
W Z
-0
-
-10
-
-11
-
-12
-
w
FeTAP
1.'
Figure 1. Molecular orbital diagram for iron tetraazaporphyrin. The lower set of orbitals is predominantly ring in character,whereas the upper set is antibondingbetween the metal and the ring and mostly metal d in character. The ?r interaction between the metal and the ring is explained in more detail in Figure 2.
formed may be expected to be mostly responsible for the overall strength of the metal-ring bond. Electronic Properties of the Primitive Fragment: Tetraazaporphyrin, FeTAP The MO diagram for FeTAP obtained from the ASED-MO theory calculations is shown in Figure 1. The alg, big, eg(n), and bzg FeTAP MO's are largely ring localized because of the large energy difference between the ring and metal 3d orbitals. The antibonding counterparts, which have large metal d coefficients, are customarily called the d,,, dyz (or d(n)), d,z, d,, and d X y orbitals. As expected, d,i-? is highly destabilized by the u mteraction with the big,ring orbital, and d, and d z are only slightly destabilized by the n in-plane and u interactions with the respective bzr and alg ring orbitals. The extent of these energy shifts can be rationalized by second-order perturbation t h e ~ r y : ~
where the xI0's are atomic or molecular orbital basis functions, H' is the interaction Hamiltonian, and the E,'s are eigenvalues of the xi's. The out-of-plane ?r interaction, only partially depicted in Figure 1, is shown in more detail in Figure 2. The empty le,(n*) TAP orbitals are close in energy to the d(.rr), so there is a strong interaction. Two degenerate sets of ring orbitals (e,(n) and 2eg(.rr*)) can also mix with the metal d(n) orbitals. As shown in Figure 2, the resulting leg(n*) M O s of FeTAP (not depicted in,Figure 1) are pushed down by the high-lying 2eK(a*)ring orbital and the bonding dn orbitals are pushed up by the low-lying ring e,(x) MOs. The result is an energy gap of less than 1 eV between the partially occupied d n and the empty leB(n*)orbitals of FeTAP. This small splitting does not mean that the mixing of the d(n) and the le,(**) TAP orbitals is weak. In fact, the back-bonding is so strong that, from the Mulliken population analysis, it is found
Figure 2. Metal d,, dyzorbitals interacting with the e8(r), le?(r*), and 2eg(n*) ring MOs of TAP. The metal to ligand back-bonding is controlled by the d,,, d,, ring LUMO energy gap (AE).
that only 1.7 out of the possible three electrons associated with these degenerate orbitals are localized on the metal. The energy level ordering of the d orbitals follows d+? > d, > dZz> d(n). The orbital occupation shown in Figure 1 is suggested by the experimental charge density measurements of FePc by Coppens et al.24 The charge a t the metal center (pM) for FeTAP was found to be +2.66, indicating a high electron acceptor character for the TAP ring. Though the lack of charge selfconsistency in the MO calculations makes the charge transfer approximate in magnitude, any structural changes which decrease this calculated number are likely to make the iron center a better reducing agent. It is precisely the aim of this study to examine the way by which ring substituents and geometry modify the charge density at the metal center. On the basis of perturbation theory arguments, a way of increasing the electronic charge density on the metal would be to increase the energy gap, AE,between the metal d(7) orbitals and the TAP LUMO. The TAP LUMO mixes with the metal d ( r ) HOMO in a way inversely proportional to the energy gap according to a first-order perturbation expre~sion.~
The greater the energy gap in the denominator, the smaller the delocalization of d(n) electron into the TAP LUMO and the less positive the metal charge. For a given metal this requires a destabilization of the TAP LUMO. This goal may be achieved (24) (a) Coppens,P.; Liang, L.; Zhu, N. J. J. Am. Chem. Soc. 1983,105, 6173. (b) Coppens, P. J . Chem. Educ. 1984,61, 761.
The Journal of Physical Chemistry, Vol. 92, No. 24, 1988 6905
Electron Donor-Acceptor Properties of Macrocycles
TABLE Ik Number of Electrons in the d, dyzOrbitals (d(x) Orbitals) and Energy Gap ( A E ) between These Orbitals and the Ring LUMO d(*) electrons AE, eV, d(*)-LUMO
/-\
/
energy gap
FeTAP FePc FeP(1) FeP(2) FeNPc 1.86 2.53 2.66 2.80 2.58 0.26 0.63 0.78 0.75 0.68 +2.66
PMa
+1.99 +1.94
+1.81
+1.93
is the charge at the metal in the Mulliken definition.
\
-8 ’
-9 .
-10
’
w Y W
’
>
8y
-11
’
-12
’
w
-I 3
Figure 3. Destabilization of the ring LUMO leg(**) of FeTAP by a butadiene fragment, an electron r donor group, leading to the phthalocyanine molecule.
by interacting the TAP LUMO fragment with a ?r donor group located a t the periphery of the ring or achieved by an electronegativity pert~rbation:~such as might be induced by replacing N by C in the bridge position of TAP. Both of these possibilities are analyzed below. Peripherally Induced LUMO Destabilization: Iron Phthalocyanine Among the many ?r electron donor groups that might be considered for inducing the desired destabilization, the butadiene moiety deserves immediate attention because the physicochemical properties of the resulting FePc have been extensively investigated. A schematic diagram for the TAP ring LUMO (RL) and the H O M O of a butadiene fragment (FH) interaction is shown in Figure 3. The increase in the metal d(?r)-LUMO energy gap ( A E ) can be rationalized by using (1). In this case
(3) where H’RL,, is the interaction for these orbitals and E , and ERL are the orbital energies. Since ERL - Em < 0, the ring LUMO is destabilized. This effect can be expected to be more pronounced upon addition of four butadiene groups to the TAP ring forming FePc. This destabilization was seen in earlier calculations.26 There the FePc le,(.*) shifted to positive values compared with FeTAP. Hence, the iron center in FePc, in contrast to FeTAP, would be expected to exhibit a weaker A interaction and less back-bonding to the ring. The M O calculations show the LUMO shifts up about 0.36 eV (Table 11). This is also reflected in the Mulliken population analysis for the d(s) orbitals; e.g., of the two d, electrons, only one is lost to the ring ligand. (25) Albright, T. A.; Burdett, J. K.; Whangbo, M.-H. In Orbirol Inreractions in Chemistry; Wiley-Interscience: New York, 1985. (26) Schaffer, A. M.; Gouterman, M.; Davidson, E. R.Theor. Chim. Acta 1973, 30, 9.
Figure 4. Assigment of the W-visible spectrum of FePc obtained in this work.
The calculated value of p~ decreases from +2.66 in FeTAP to 1.99 in FePc. The calculated molecular orbital diagram for FePc is similar to that for FeTAP. The energy level ordering of the metal orbitals, d+2 >, dxh> d s > d ( r ) is in agreement with Gouterman’s calculation, except that the dv splitting reported by these workers was not observed in the present study. The assumed electronic ground-state configuration (d,)2(d2)’(d(~))3 comes from Coppens’ charge density determinations using X-ray data:4 which suggests double occupation for dXy This occupation was proposed earlier by Gouterman et a1.26 The transitions from the predominantly ring al,(?r) orbital, and the lower lying ring azu(a)orbital, to the ring LUMO le (.*) are responsible for the characteristic visible (Q) and Soret (B) absorption bands. The values obtained in the calculations are close to those obtained experimentally, 620 nm (630) for Q and 440 nm (420) for B, where experimental values are in parentheses. These and other 7 ~transitions * are indicated in Figure 4. The le,(.rr*) orbital at -9.5 eV was shown in Figure 2 for FeTAP.
+
Inner Ring Induced LUMO Destabilization: Iron Porphyrin Two different structures were used in order to assess an electronegativity perturbation on the chelate fragment. In the first case, the bridge nitrogens in FeTAP were replaced by carbon atoms without varying the Fe-N bond length, a geometry that preserves the coordination ligand hole size.. The molecular orbital diagram obtained for this compound, denoted hereafter as FeP( l), yielded the same M O energy ordering as that for FePc and FeTAP. The magnitudes of pM and the metal d(?r)-LUMO le,(**) energy gap were found to be larger than for FeTAP and close to those obtained for FePc (Table 11). The charge and gap increase because carbon, which is less electronegative than nitrogen and has a smaller 2p ionization potential, is in a position where the le,(**) MOs have a major contribution. The ring LUMO destabilization caused by carbon substitution at the meso position of FeTAP is slightly larger than that introduced by the butadiene fragments examined earlier (Table 11).
Fierro et al.
6906 The Journal of Physical Chemistry, Vol. 92, No. 24, 1988
The a2u(7r) orbital also has contributions at the bridge atomic positions, and hence the nitrogen-carbon substitution reduces its stability. The second species examined was iron tetraphenylporphyrin (FeTPP)27with the four phenyl groups at the periphery of the molecule omitted. This structure has a slightly larger coordination hole than FeP(1) and will be referred to as FeP(2). For this, the calculations gave increased electron population in the d(a) orbitals but practically no change in AE (see Table 11). This means that the longer Fe-N bonds of porphyrins (by 0.045 A) decrease the back-bonding to the ring. This effect can be understood from second-order perturbation theory. Specifically, an increase in the Fe-N bond length decreases the magnitude of the numerator in (2). Since AE in the denominator remains practically unchanged, this causes a decrease in the metal to ring donation. The metal character of the d(a) orbitals is small, 52% in these calculations and 37% and 42% in previous Xaz8 and EHT16 calculations, respectively. The large ring contribution to the d(a) compared with the dx2+ (63% metal character), dZ2(99%), and d, (82%) orbitals found in this work may be associated with the strong ring mixing found in the case of FeTAP (see Figure 2). Although the relative energies of these orbitals are in agreement with the extended Hiickel results,16 they are at variance with the recent calculations based on the X a multiple-scattering methodz8 where the alu(a),aZu(r)orbitals of FeP(2) were above the metal d levels. The weaker metal to ring back-donation may explain the relative ease which Fe(I1) porphyrins are oxidized. The iron in most of these oxidized species is five-coordinate and displaced by about 0.45 A from the molecular plane. At this distance Coulombic repulsion between the axial ligand and the ring is reduced, and this should enhance the overall stability of the complex (see ref 15). The present calculations indicate that Fe"Pc has 0.063 eV less stability than Fe"P(2) for a displacement of the iron center of about 0.5 A away from the molecular plane. This factor, in addition to the more difficult metal oxidation due to the increased back-bonding, might play a role in preventing the formation of stable five-coordinate FeInPc species. Indeed, many five-coordinate Fe"'P complexes have been characterized, but no Fe"'Pc analogues have been unambiguously identified. The Fe"Pc can, however, be stabilized by axially coordinated Lewis bases such as pyridine or dimethyl sulfoxidez9 (vide infra). Sabelli and Melendres14 have proposed that the d, orbital in transition-metal macrocycles is destabilized as a consequence of u antibonding with the bridge atoms and that the destabilization should be larger in phthalocyanines than prophyrins. The present calculations indeed show that the d, orbital experiences greater destabilization (by 0.29 eV) in FePc than in FeP(2), and this is actually caused by an in-plane ?r antibonding interaction involving the ciosest nitrogens as depicted in 6 and Figure 1. This, however, appears to induce a rather minor effect on the metal charge.
I
/
Figure 5. Structural progression from the basic fragment TAP to phthalocyanine and naphthalocyanine by the addition of butadiene fragments at the periphery of the molecule.
TABLE IIk
Same as Table I1 for Other Species
FeTfPc FePz(1) FePz(I1) FeTPP 2.3 2.94 d ( r ) electrons 2.58 2.35 0.46 0.86 AE, eV, d(?r)-LUMO 0.66 0.52 energy gap
+1.95
PM
+2.17
+2.22
+1.66
deactivated porphyrin 2.73 0.64
+1.90
attached to each of the isoindole units. The ring LUMO in FePc, as pointed out earlier, is mainly localized at the bridge and indole nitrogens. Since this specific peripheral substitution on Pc does not involve directly either of these two ring positions, the effects on pm and AE might be expected to be small. This was indeed found in the calculations (see Table 11). A very small change in pMrrelative to FePc, was also obtained in the case of iron tetrafluorophthalocyanine (FeTfPc; see Table 111). Indole Ring Substitution The electronic properties of two compounds described in the literature3z in which one of the carbon atoms in each of the isoindole groups in FePc is replaced by nitrogen, 7,were also
7a
7b
examined. In order to compare with our previous calculations, the geometric parameters of FePz(1) (7a) and FePz(I1) (7b) were assumed to be the same as those of FePc. Table I11 shows that the d ( r ) electron populations for FePz(1) Further Ring Modifications: Peripherally Substituted Iron and FePz(I1) are slightly lower than in FePc. In addition, the Phthalocyanine metal to ligand back-bonding is greater for FePz(II), where pM = +2.22, than for FePc, where pw = +1.99. It has been suggested by several authors30 that an increase in In accordance with the arguments put forward in the case of the size of the resonant structure of the macrocycle should result FePc, the replacement of one of the carbon atoms in the butadiene in a higher electron density at the iron center and thus give rise fragment by nitrogen leads to a decrease in the energy of the to enhancement in the electrocatalytic activity for Oz reduction. Examples of such extended x systems are naphthalo~yanines~~ fragment HOMO and consequently to an increase in the HOMO-LUMO energy gap (see Figure 3). This effect is more and certain polymeric forms of FePc3' pronounced in the case of the FePz(I1) isomer because the carbon Iron naphthalocyanine, FeNPc, as shown schematically in atom replaced by nitrogen had larger coefficients in the fragment Figure 5, may be regarded as FePc with butadiene fragments HOMO. Thus the AE in FePz(I1) was smaller than that in FePz(1) (Table 111). (27) (a) Collman, J. P.; Hoard, J. L.; Kim, N.; Lang, G.; Reed, C. A. J . Am. Chem. SOC.1975, 97, 2676. (b) Lee, K. L.; Sabelli, N. H.; LeBreton, P. R. J. Phys. Chem. 1982,86, 3926. (28) Sontum, S. F.;Case, D. A. J . Am. Chem. SOC.1985, 107, 4013. (29) (a) Lever, A. B. P.; Wilshire, J. P. Inorg. Chem. 1978, 17, 1145. (b) Lever, A. B. P.; Wilshire, J. P. Can. J. Chem. 1976, 54, 2514. (c) Lever, A. B. P. Ado. Chem. Ser. 1979, 173, 237. (30) Magner, G.; Savy, M.; Scarbeck, G.; Riga, J.; Verbist, J. J. J . Electrochem. SOC.1981, 128, 1674. (31) (a) Savy, M.; Andro, P.; Bernard, C.; Magner, G . Electrochim. Acta 1973,18, 191. (b) Savy, M.; Andro, P.; Bernard, C. Electrochim. Acta 1974, 19, 403.
Meso-Substituted Iron Porphyrins The hydrogens in the meso positions of FeP(2) can be replaced by other functional groups, providing a means of modifying the macrocycle ring electron acceptor-donor capabilities. Phenyl meso ligands are particularly attractive since they can be further (32) Smith, T.D.; Lirorness, J.; Taylor, H. J. Chem. SOC.,Dalton Trans. 1983, 1391.
Electron Donor-Acceptor Properties of Macrocycles functionalized by a variety of substituents. The effects of attaching four phenyl groups to the meso position of FeP(2) have been examined by using the atomic coordinates reported by Collman et al.” for FeTPP. Two modifications of the actual structure were introduced to simplify the M O analysis. First, the porphyrin ring was deformed so as to acquire Dghsymmetry, and second, the phenyl groups were assumed to be perpendicular to the porphyrin plane instead of at the 80° angle found experimentally. This nearly normal orientation with respect to the molecular plane arises from steric factors involving the ring pyrrole moieties. As shown in Tables I and 11, the destabilization of the ring LUMO induced by this substitution is larger by 0.1 1 eV than that in FeP(2), leading to a decrease in the metal to ring back-bonding. This in turn increases the electron population of the d(?r) orbitals, yielding most probably a molecule with a stronger reducing power. A more pronounced effect of the phenyl substituents could be achieved by increasing the distance between the porphyrin ring and the phenyl groups so as to avoid steric hindrance and thus extended the ?r resonant system. This can be realized, at least theoretically, by adding acetylene linkages between the moieties, as shown in 8. Calculations involving this hypothetical molecule
‘2-0 1’
8
were performed by using standard C-C triple bond distances. As shown in Table 111, this modification, denoted as deactivated porphyrin, gave rise to a decrease in Ai3 of about 0.22 eV with respect to the FeTPP and thus to an increase in the metal to ring back-bonding. The latter is reflected in the more positive pM, which is due mostly to the decrease in the population of the d(?r) orbitals.
Axially Coordinated FePc Another means of perturbing the electronic density of the iron center is ligand coordination to the axial position of the macrocycle. Preliminary calculations involving Lewis bases axially bound to Fe in FePc have shown that the u component of the donation interaction shifts the d,z orbital toward higher energies. If the energy gap generated between the d z and the d(?r) orbitals is large, the electron may drop to the lower energy level bringing about a low spin configuration (S = 0). This has indeed been observed experimentally for a variety of complexes of the form FePcL,, where L represents a Lewis base: they are ESR silent. Furthermore, the r-donation destabilization of the d ( r ) orbitals, although not as pronounced as for the d i orbital, appears to be strong enough to increase the back-bonding from the metal to the ring LUMO, which is expected to decrease the reducing capability of the metal center. It is to be emphasized that this decrease in AE is attributed to a perturbation of the metal orbitals involved in the back-bonding rather than to the ring acceptor orbital destabilization examined previously. In some cases the value of pM was found to increase up to 20% over that of FePc for model axially coordinated species, a result of the stronger back-bonding. On the basis of these considerations, FePc with an axially coordinated Lewis base can be regarded as a modified macrocycle with less d(n) and d,i occupation than FePc. The smaller value of pM for a singly axially bound compound might result in a weaker interaction with an electron acceptor ligand (such as 0 2 ) coordinated in the remaining axial position.
The Journal of Physical Chemistry, Vol. 92, No. 24, 1988 6907 The similarities between the axially coordinated FePc and FeTAP suggest that the extent of back-bonding in both molecules is comparable and stronger than in FePc without any axial ligand. In the case of FePc, the back-bonding is enhanced by the presence of the axial ligand, thus stabilizing the iron(I1) complex without modifying the macrocycle ring structure. Furthermore, for an axially coordinated species the d z orbital is empty, and since the d(r) orbitals are less populated, this represents favorable situation for a basic ligand to interact at the 6-position of FePc. An analogous macrocycle with a metal center having more d electrons such as CoPc may, with axial ligation, undergo a more facile oxidation than FePc since the electron in the destabilized dZzorbital may not find (within a certain energy range) a lower available empty orbital in the same molecule (as in FePc). This appears to explain the experimental observations reported by Weber and Bush.33 Finally, it is mentioned that additional calculations have indicated that the axial coordination of O2to the iron center in FePc gives rise to the formation of an unstable adduct which exhibited no barrier for dissociation of dioxygen. While ASED-MO calculations may underestimate the energy barrier to cleavage of coordinated 02, this reaction could be involved in the formation of a p-oxo species or under certain conditions to the destruction of the macrocycle ring. In fact, Ercolani et al.34 have reported that the continuous bubbling of O2into a very dilute aqueous solution of FeTsPc leads to the complete disappearance of the characteristic blue color and the formation of a white precipitate. Summary The electronic properties of iron macrocycles and in particular the charge density at the metal center, pM,are controlled by the energy difference between the d(?r) metal orbitals and the lowest unoccupied ring orbital LUMO, AE. The strength of such metal-ring interactions has been assessed with iron tetraazaporphyrin (FeTAP) as a primitive fragment. On this basis, the properties of a variety of iron macrocycles have been rationalized. Significant changes in the value of AE are observed upon directly perturbing the 18 ?r electrons of FeTAP by four butadiene groups or by simply replacing the nitrogen bridge atoms by carbon. This leads to phthalocyanine (Pc) and porphyrin-type compounds (P( l)), respectively. A further structural modification of FeP( 1) to yield the correct Fe-N bond distance for actual porphyrins (FeP(2)) decreases the degree of back-bonding between the metal and the ring as a result of a weaker ?r-orbital overlap without changing the value of AE. Further modifications involving either peripheral or inner ring substitutions to FePc (naphthalocyanines and porphyrazines) were found to yield small changes both in AE and pM. In the case of porphyrins, the attachment of substituents to the meso positions provides a means to further modify the values of such parameters. This indicates that only a direct perturbation to the 18 ?r electron system results in an appreciable change in the electronic structure of the primitive fragment (FeTAP).
Acknowledgment. Support for this work has been provided by the Department of Energy through a subcontract with the Lawrence Berkeley Laboratory, the Office of Naval Research, the Gas Research Institute, the Diamond Shamrock Corp., and the International Business Machines Corp. through a Faculty Development Award to D.A.S. We express our appreciation to Prof. E. B. Yeager for helpful discussions. (33) Waver, J. H.; Bush, D.H.Inorg. Chem. 1965, 4, 412. (34) Ercolani, C.; Gardini, M.; Monacelli, F.; Pennesi, G.;Rossi, G . Inorg. Chem. 1983, 22, 2584.
