Synthesis and Characterization of a Tripyrrane-Copper(II) Complex

Full text of DOI:10.1021/ic9606789

6636
Inorg. Chem. 1996, 35, 6636-6637
Synthesis and Characterization of a Tripyrrane-Copper(II) Complex
Jonathan L. Sessler,* Andreas Gebauer, Vladim´ır Kra´l, and Vincent Lynch
Department of Chemistry and Biochemistry, University of Texas at Austin, Austin, Texas 78712
ReceiVed June 7, 1996
To date, the study of linear oligopyrroles, systems containing
several pyrrolic moieties linked by methylene or methene
fragments, has been largely focused on the chemistry of naturally
occurring bile pigments.1 These pigments, which contain four
pyrrolic fragments arranged in a linear fashion, have so far been
used to stabilize complexes of Fe(II), Co(II), Mn(II), and
Cu(II).2 Bilirubin and related materials are thus emerging as
biogenic ligands of considerable interest. However, the very
success of these tetrapyrrolic chelands serves to raise an obvious
question as to whether linear polypyrrolic systems containing
a greater or lesser number of pyrroles might not also be
employable as metal-complexing agents. In this communication
we describe the synthesis of a Cu(II) complex (2) derived from
an oxidized tripyrrane ligand.
Tripyrranes (e.g., 1)3 are one of the better studied oligo-
Figure 1. View of complex 2, C42H45N3O7Cu‚0.5 CHCl3, showing
pyrroles. They have been used extensively as building blocks
in the synthesis of porphyrins4 and expanded porphyrins,5
including sapphyrin,6 pentaphyrin,7 hexaphyrin,8 and texa-
phyrin.9 When incorporated into the latter macrocycles, the
tripyrranes and their 4-electron-oxidized congeners, the tri-
pyrrodimethenes, display a rich metal-binding chemistry. How-
ever, it remains an open question at present whether these non-
natural “bile pigment analogues” can themselves act as ligands
absent any kind of macrocyclic effect. It is known, however
that dipyrromethenes can function in this regard.10
the atom-labeling scheme. Thermal ellipsoids are scaled to the 30%
probability level. Hydrogen atoms have been omitted for clarity.
tripyrrane ester 1 to copper(II) acetate hydrate under oxidizing
conditions (acetonitrile, trace triethylamine, reflux) leads to an
apparent reaction (Scheme 1). Subjecting the resulting product-
(s) to successive column chromatographic purifications (silica
gel, dichloromethane/1% methanol eluent) was found to give
rise to only one blue-green fraction. After removal of solvent,
this fraction gave a green solid that proved paramagnetic as
judged by NMR spectroscopy. Mass spectrometric analysis of
this product revealed a M+ peak that was 14 mass units heavier
than expected for a copper(II) tripyrrodimethene complex. This
led us to consider the possibility that water or some other simple
fragment was incorporated into the tripyrrane fragment during
the course of metal complexation. However, these data did not
tell us how or where this putative fragment was “inserted”.
Fortunately, the above complex proved stable as a dichlo-
romethane solution when either stored at room temperature in
the presence of air for several months or exposed briefly to
aqueous solutions of pH 2-14.11 It thus proved possible to
crystallize it from chloroform/pentane. This, in turn, made it
possible to resolve the “mystery” Via X-ray diffraction analysis.
The X-ray crystal structure of the tripyrrane-Cu(II) complex
(2) is shown in Figure 1.12 It reveals that, in the course of metal
chelation, the tripyrrane ligand is oxidized at one of the two
bridging methylene positions to form a bridging methene and
at the other to form an exocyclic keto group. This surprising
oxidation process converts the starting tripyrrane into a species
with two acidic protons. It thus generates a ligand that is able
to coordinate Cu(II) in the form of an overall neutral complex.
As revealed by the X-ray structure, the coordination environ-
ment of the Cu(II) center in 2 is a distorted square plane. The
To test the above possibility we sought to prepare a
representative tripyrrane-metal complex. After considerable
experimentation, we found that exposure of the dibenzyl
(1) Falk, H. The Chemistry of Linear Oligopyrroles and Bile Pigments;
Springer Verlag: Vienna, 1989.
(2) (a) Balch, A. L.; Mazzanti, M.; Noll, B. C.; Olmstead, M. M. J. Am.
Chem. Soc. 1994, 116, 9114-9122. (b) Balch, A. L.; Mazzanti, M.;
Noll, B. C.; Olmstead, M. M. J. Am. Chem. Soc. 1993, 115, 12206-
12207.
(3) Sessler, J. L.; Johnson, M. R.; Lynch, V. J. Org. Chem. 1987, 52,
4394-4397.
(4) (a) Boudif, A.; Momenteau, M. J. Chem. Soc. Chem. Commun. 1994,
2069-2070. (b) Nguyen, L. T.; Senge, M. O.; Smith, K. M. J. Org.
Chem. 1996, 61, 998-1003. (c) Sessler, J. L.; Genge, J. W.; Sansom,
P. I.; Urbach, A. Synlett 1996, 2, 187-188. (d) Berlin, K.; Breitmaier,
E. Angew. Chem., Int. Ed. Engl. 1994, 33, 1246-1247. (e) Lin, Y.;
Lash, T. D. Tetrahedron Lett. 1995, 36, 9441-9444.
(5) (a) Dolphin, D., The Porphyrins Vols. I and IIB, Academic Press, New
York, 1978. (b) Sessler, J. L.; Burrell, A. Top. Curr. Chem. 1991,
161, 178-273.
(6) (a) Broadhurst, M. J.; Grigg, R.; Johnson, A. W. J. Chem. Soc., Perkin
I 1972, 2111-2116. (b) Bauer, V. J.; Clive, D. L. J.; Dolphin, D.;
Paine, J. B., III; Harris, F. L.; King, M. M.; Loder, J.; Wang, S.-W.
C.; Woodward, R. B. J. Am. Chem. Soc. 1983, 103, 6429-6436. (c)
Sessler, J. L.; Cyr, M.; Burrell, A. K. Tetrahedron 1992, 48, 9661-
9672 and references therein.
(7) (a) Rexhausen, H.; Gossauer, A. J. Chem. Soc., Chem. Commun. 1983,
275. (b) Gossauer, A. Chimia 1983, 37, 341-342. (c) Gossauer, A.
Chimia 1984, 38, 45-46.
(8) Gossauer, A. Bull. Soc. Chim. Belg. 1983, 92, 793-795.
(9) (a) Sessler, J. L.; Mody, T. D.; Hemmi, G. W.; Lynch, V. Inorg. Chem.
1993, 32, 3175-3187. (b) Sessler, J. L.; Hemmi, G. H.; Mody, T. D.;
Murai, T.; Burrell, A.; Young, S. W. Acc. Chem. Res. 1994, 27, 43-
50.
(10) For examples see: (a) Corwin, A. H.; Sydow, V. L. J. Am. Chem.
Soc. 1953, 75, 4484-4486. (b)Murakami, Y.; Sakata, K. Inorg. Chim.
Acta 1968, 2, 273-279. (c) Fritschi, H.; Leutenegger, U.; Siegmann,
K.; Pfaltz, A.; Keller, W.; Kratky, C. HelV. Chim. Acta 1988, 71,
1541-1552.
(11) The stability of complex 2 was determined kinetically by monitoring
the change in the Q-band absorbance at 608 nm as a function of pH
using a Beckmann DU-7 spectrophotometer. This was done by
exposing complex 2, dissolved in dichloromethane, to aqueous
solutions of a given pH between 1 and 14.
(12) Crystal data and details of the data collection and structure refinement
are given in the Supporting Information.
(13) (a) Hathaway, B. J. Essays Chem. 1971, 2, 61-92. (b) Cotton, F. A.;
Wilkinson, G. AdVanced Inorganic Chemistry, 4th ed.; Wiley Inter-
sience: New York, 1980; p 813.
S0020-1669(96)00678-7 CCC: $12.00 © 1996 American Chemical Society

Communications
Inorganic Chemistry, Vol. 35, No. 23, 1996 6637
nm) as compared to the corresponding bands of the cobalt and
zinc complexes (λmax ≈ 540 nm).
The electrochemical properties of the copper(II) complex 2
were probed using cyclic voltammetry.14 When the complex
is scanned between +400 and -1000 mV, a reversible Cu(II)/
Cu(I) wave (E1/2 ) 480 mV vs Ag/AgCl) is observed (Sup-
porting Information). On the other hand, when the cyclovol-
tammogram is recorded between +2 and -2 V, a second
reduction wave is observed (Supporting Information) that is
tentatively assigned to a ligand-centered reduction process.
Under these scanning conditions, the Cu(II)/Cu(I) couple
becomes irreversible. This leads us to propose that once the
ligand is reduced, some irreversible, presumably dissociative,
process takes place such that the copper(I) cannot be oxidized
back to copper(II).
The EPR spectrum of the copper(II)-tripyrrane complex 2
at 115 K was also recorded (Supporting Information).15 From
these measurements, an A| value of about 150 G was determined
for complex 2. While the observed signals are very broad
(making an exact determination of A| difficult at this temper-
ature), the fact that the A| value for 2 is on the low side of what
is generally seen for normal Cu(II) complexes (for which A|
values of 140-200 G are generally recorded16) leads us to
suggest a degree of interaction between the copper(II) ion center
and the tripyrrane ligand. Such a putative delocalization is
consistent with what is observed by other spectroscopic means
(Vide supra). However, simple geometry-based explanations
for this low value (i.e., distortion from pure square planarity)
are also possible.
Figure 2. UV/vis spectra of the Cu(II)-, Co(II)-, and Zn(II)-
tripyrrane complexes 2, 3, and 4 in dichloromethane.
Scheme 1
Scheme 2
The present results serve to illustrate that “bile pigment
analogues” containing but three pyrroles can be used as acyclic
ligands for copper(II) and, seemingly, other transition metal
cations. On this basis, we are inclined to propose that a wide
range of oligopyrrolic species, including those containing five
and six pyrrolic subunits,17 could be used to support the
formation of interesting metal complexes.
copper atom is bound covalently to N1 and N3 at bond lengths
of 1.878(5) and 1.909(5) Å, respectively, while N2 (Cu-N2
1.938(6) Å) and O2 (Cu-O2 2.156(5) Å) of the R-ester
functionality are used to fill the coordination sphere of the
copper atom. The angles involving N1, N2, N3, and Cu (N1-
Cu-N3 173.8(2)°, N1-Cu-N2 90.3(2)°, N2-Cu-N3
94.2(2)°) resemble, with only small deviations, those expected
for a square planar complex.13 The angles involving O2, Cu,
and N1-N3 (N1-Cu-O2 79.4(2)°, N2-Cu-O2 157.4(2)°,
N3-Cu-O2 97.8(2)°), on the other hand, are distorted sub-
stantially from those expected for a metal bound in an idealized
square planar ligand environment. This deviation is presumably
a reflection of the steric requirements of the R-ester; this group
must be bent into the main chelation plane in order to interact
with the copper atom.
To explore whether such tripyrrane species can also act as
ligands for other metal cations, two additional metal complexes
were prepared. This was done by first demetalating the copper
complex 2 with concentrated HCl and then treating the reaction
mixture with either cobalt acetate or zinc acetate. The resulting
complexes (3 and 4) were purified by chromatography and
characterized by mass spectrometry and UV/vis spectroscopy.
Figure 2 shows an overlay of the UV/vis spectra of the
cobalt-, copper-, and zinc-tripyrrane complexes 2, 3, and 4.
Both the cobalt- and zinc-oxotripyrromethene complexes
display similar spectra, with the major difference being the
higher extinction coefficients of the cobalt(II) species. The
copper(II) complex, on the other hand, displays a markedly
different spectrum. It shows a greater molar absorptivity and a
stronger Q-type band. This latter is red-shifted (to about 608
Acknowledgment. This work was supported by a grant from
the National Science Foundation (CHE 9122161) to J.L.S.
Supporting Information Available: Text giving synthetic and
X-ray experimental procedures, tables of crystallographic data, atomic
thermal parameters, fractional coordinates, and bond lengths and angles
for the tripyrrane-Cu(II) complex 2, an EPR spectrum of complex 2,
cyclovoltammograms of complex 2 scanned between +0.4 and -1 V
and between +2 and -2 V, and additional structural diagrams of 2
(24 pages). Ordering information is given on any current masthead
page.
IC9606789
(14) Cyclic voltammetry was carried out at room temperature (25 ( 2 °C)
under dry argon using a Bioanalytic Systems Inc. (BAS) CV-50W
Version 2 MF 9093 voltammetric analyzer. Dry acetonitrile was the
solvent, 0.1 M [CH3(CH2)3]4NPF6 the electrolyte, a platinum disk the
working electrode (1.6 mm diameter), and a platinum wire the auxiliary
electrode. A Ag/AgCl couple, separated from the bulk solution by
means of a porous Vycor plug, was used as the reference electrode.
(15) Recorded on a Bruker ESP 300 EPR spectrometer interfaced to an
IBM ESP 1600 computer.
(16) Solomon, E. I.; Lowery, M. D. In The Chemistry of the Copper and
Zinc Triads, Welch, A. J., Chapman, S. K., Eds.; Royal Society of
Chemistry: London, 1993; see also references therein.
(17) Sessler, J. L.; Weghorn, S. J.; Lynch, V.; Fransson, K. J. Chem. Soc.,
Chem. Commun. 1994, 1289-1290.