Mixed valency and metal–metal quadruple bonds

Coordination Chemistry Reviews 257 (2013) 1576–1583

Contents lists available at SciVerse ScienceDirect

Coordination Chemistry Reviews journal homepage: www.elsevier.com/locate/ccr

Review

Mixed valency and metal–metal quadruple bonds Malcolm H. Chisholm Department of Chemistry, The Ohio State University, 100 West 18th Avenue, Columbus, OH 43210-1185 USA

Contents 1. 2. 3. 4. 5. 6. 7.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Electrochemical studies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Origins of electronic coupling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Electronic absorption spectral features . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Other spectroscopic techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Mixed valence excited states . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Concluding remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Acknowledgements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

a r t i c l e

i n f o

Article history: Received 28 June 2012 Received in revised form 31 October 2012 Accepted 31 October 2012 Available online 7 December 2012

1576 1577 1577 1578 1580 1580 1582 1582 1583

a b s t r a c t Linking organic conjugated systems with MM quadruple bonds (M = Mo or W) allows for the study of mixed valence metal and organic centered radicals in their ground and photo-excited states. A variety of spectroscopic studies have been employed to determine the nature of the charge distribution and the classification as Class II or III on the Robin and Day scheme. © 2012 Elsevier B.V. All rights reserved.

Keywords: Molybdenum Tungsten Quadruple bonds Spectroscopy

1. Introduction In 1989 I initiated a project to look into the properties of MM multiple bonds in ordered assemblies [1]. The well-defined paddle-wheel like structure of the dimetal tetracarboxylates of molybdenum and tungsten provided an excellent building block and entry point into this field of investigation [2]. Their ability to bind ligands, albeit weakly along the MM axis, leads them to self-associate in the solid state and form laddered structures. This self-assembling allowed various n-alkanoates of dimolybdenum to be studied in their thermotropic, magnetically aligned mesophases [3,4]. Early attempts to prepare polymeric chains incorporating metal–metal quadruply bonded units involved the reactions between the M2 (O2 CR)4 homoleptic compounds and dicarboxylic acids or from reactions involving salt metathesis with the [Mo2 (O2 CBut )2 (CH3 CN)n ]2+ cation [5,6]. These reactions turn

E-mail address: [email protected] 0010-8545/$ – see front matter © 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.ccr.2012.10.021

out to be much more complex than they would appear on paper and lower molecular weight oligomers and cyclic products such as molecular loops, triangles, and squares are encountered with various bridging groups [7,8]. This line of research continues in my lab though over the course of the years many excursions have been made, one of which is the study of mixed valency in covalently linked MM quadruply bonded units. In retrospect these systems are ideally suited for the study of mixed valence as oxidation of the quadruple bond leads to the removal of an electron from the ␦ orbital to yield the MM bonding configuration ␴2 ␲4 ␦1 . Crystallographically it is easy to distinguish the difference between a MM bond of order 4 with one of order 3.5, since this oxidation is ˚ This accompanied by a lengthening of the MM distance by ∼0.04 A. is readily detected with current crystallographic methods which most often can determine the distances between these 4d and 5d metals to within ±0.001 A˚ [2]. Cotton, Murillo, and coworkers have prepared numerous examples of crystallographically characterized Mo2 linked units supported by formamidinates in their neutral and oxidized forms and thus detected valence trapped or delocalized ions crystallographically [9]. In fully delocalized linked M2

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

containing mixed valence ions the MM bond order is 3.75 which can similarly be distinguished from M2 centers with bond order 3.5 and 4. However, despite the power of single crystal X-ray studies, there can still be problems associated with disorder and for an oxidized ion the presence of the counter anion in the solid-state can lead to charge pinning. Consequently, solution studies are really invaluable. 2. Electrochemical studies Cyclic voltammetry and differential pulse voltammetry are particularly useful screening methods in the study of mixed valence as was pioneered by Richardson and Taube [10]. For linked quadruply bonded compounds the separation between the first and second redox potentials, E in mV, relates the stability of the mixed valence ion in equilibrium with the neutral and doubly oxidized species. The comproportionation constant, Kc is equal to e(E1/2 /25.69) under standard conditions. In this way it was very soon apparent to us that the nature of the bridging ligand is critically important in coupling the two dinuclear centers. For example, the 2,7-oxynaphthyridine ligand and 1,8-anthracenyldicarboxylate are stereochemically correspondent, as shown in I and II below, in bringing the M2 · · ·M2 distance to within 3 A˚ separation yet the E1/2 value of the former was much greater than the latter [6].

However, the oxalate bridge, which separates the two M2 ˚ is even more effective as seen for the comcenters by ∼6 A, plex [(But CO2 )3 W2 ]2 (␮-O2 CCO2 ). Selected electrochemical data are given in Table 1. The electrochemical data, though easy to obtain, do not clearly indicate whether an ion is fully delocalized as in Class III on the Robin and Day classification scheme [11]. The data reflect the thermodynamic properties of the ions present and as such are very dependent on the nature of the solvent and the conditions. This mater has been well documented by Geiger in comparing the use of non-coordinating ions such as B(Arf )4 , where Arf = 3.5-(CF3 )2 C6 H3 [12]. However, what is readily apparent from the electrochemical data is that the magnitude of the electronic coupling is always greater for tungsten relative to molybdenum containing complexes. Given that the two elements are of almost identical size with similar solution energies, the greater coupling can be directly correlated with orbital interactions. 3. Origins of electronic coupling The electronic coupling arises from the interaction of the M2 ␦ orbitals and the bridging ligand ␲-system. The 5d element tungsten has an ␦ orbital oxidation potential roughly 0.5 V less than that of its 4d molybdenum counterpart. This, together with the greater 5d overlap with the bridging ligand orbitals, emphasizes the importance of the back-bonding in the electronic coupling of the two M2 centers. For carboxylate bridging ligands one of the ␲* bridging

1577

combinations will find a match with either the symmetric or anti-symmetric combination of the two M2 ␦ orbitals. Similarly for a single M2 center with two trans ␲-accepting carboxylate ligands, the out-of-phase ligand ␲* combination has a symmetry match with the M2 ␦ orbital. These complementary coupling of M2 centers and ligands are shown pictorially in Fig. 1. Here we are emphasizing the importance of the M2 ␦ to bridge/ligand ␲* orbital interactions. Other interactions are also possible by symmetry, namely the filled CO2 ␲ orbitals with the M2 ␦ but based on energy considerations this is less important. The replacement of the carboxylate oxygen by sulfur or an NR group raises the energy of the filled ␲ bridge/ligand orbital and thus further facilitates M2 bridge/ligand bonding. For the bridged M2 complexes the magnitude of the splitting of the two ␦ orbitals represents a measure of the electronic coupling. Similarly for the two trans ␲-accepting ligands, the energy splitting between the two ligand ␲* systems is an indicator of how the two ligands are coupled through the metal center. Just as the M2 bridged complexes show two reversible oxidation waves, the mixed valence anions show two reduction waves. From the simplified frontier MO diagrams depicted in Fig. 1, it becomes apparent that oxidation of the bridged M2 complexes and reduction of trans-ligated M2 complexes will lead to mixed valence metal centered cations and ligand centered anions,

respectively. Furthermore for the cation, a low energy electronic transition is expected involving the promotion of an electron from the HOMO − 1 to the singly occupied HOMO. For the anion, the promotion of an electron from the singly occupied LUMO to the LUMO + 1 would yield the same. These represent the solvent independent “IVCT” bands or the charge resonance bands of Class III ions on the Robin and Day classification scheme [11]. These qualitative expectations are nicely supported by calculations based on density functional theory and time-dependent density functional theory that also yield predictions of relative energies and oscillator strengths for the electronic transitions. In general these predictions are really in line with experimental observations for the molybdenum complexes but in more qualitative agreement for the related tungsten complexes for which both spin–orbit coupling and relativistic effects are more important. If the magnitude of the electronic coupling is not sufficiently large then the MO picture does not pertain and a solvent dependent higher energy and broader transition will be observed of the type and form predicted by Hush for Class II mixed valence ions [13]. In addition to these low energy transitions that occur in the infrared, IR or near-IR both ions will also have a metal to bridge/ligand charge transfer transition at higher energy. These matters have been discussed in detail elsewhere and it is sufficient here to exemplify these trends as observed by electronic absorption spectroscopy [14,15].

1578

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

Table 1 Comparison of data for selected bridged MM quadruply bonded complexes and their mixed valence ions. Bridge (2-)

Metal

1

MLCT (nm)

Kc

HAB (cm−1 )

References

Mo W

460 800

5.4 × 10 1.3 × 1012

2000 2980

[6,22] [6,22]

Mo

550

∼4

∼400

[6,23]

W

727

5.1 × 102

1500

[6,23]

Mo

490

13

–

[6,23]

W

900

6.6 × 104

1500

[6,23]

Mo

645

1.3 × 103

–

[32]

W

777

5.7 × 108

2160

[32]

Mo W

700 1140

78 8.2 × 107

380 1540

[19] [19]

Mo W

450 704

1.7 × 107 4.5 × 1010

– 1900

[33] [33]

4. Electronic absorption spectral features The electronic absorption spectra of the mixed valence oxalate bridged cations [(But CO2 )3 MM]2 (␮-O2 CCO2 )+ in THF are shown in Fig. 2 in the NIR region. The relatively narrow and symmetrically shaped absorptions for the complexes where MM = W2 or MoW are readily assignable to the fully delocalized Class III mixed valence ions where the magnitude of the electronic coupling, HAB is just one half the value of ¯ (cm−1 ) for the absorption maximum. The absorption profile for the mixed valence Mo2 containing ion is notably different. Its absorption maximum is at lower energy and the absorption is notably asymmetric. It has a relatively sharp onset at low energy but has a distinct tailing to higher energy. This is exactly as predicted for an ion on the Class II/III boarder [16]. The absorption profile arises from the fact that the electronic coupling results in a relatively flat ground state potential energy surface. As the coupling between the metal centers increases in the order MM = Mo2 < MoW < W2 the ground state potential energy

4

surface becomes more parabolic and nested with respect to the higher energy surface thus yielding a sharper more symmetric absorption. In general, we have found the radical anions of the form trans-[L-M2 -L]− are much less stable and difficult to study due to their limited kinetic persistence. As noted earlier, evidence of the coupling of the two ligand ␲ systems is indicated by the appearance of two ligand reduction waves. For the compounds trans-M2 (Ti PB)2 [O2 C-p-C5 H4 N B(Arf)3 ]2 , which contain the 4isonicotinate ligands bonded to the tris(pentafluorophenyl)boron Lewis acid, the singly reduced species is sufficiently persistent to be studied by spectroscopic techniques [17]. The radical anion has an EPR signal at g ∼ 2.0 lacking any hyperfine coupling to the M2 center and is typical of an organic radical. In the NIR low energy electronic transitions are seen. These are shown in Fig. 3. These again can be classified as Class III mixed valence ions where HAB is greater for the W2 complex. The Mo2 complex also shows the asymmetry in its absorption typical of a mixed valence ion close to the Class II/III border [16].

Fig. 1. Schematic frontier molecular orbital diagram showing the relationship between M2 -bridge-M2 and trans L-M2 -L complexes in forming mixed valence ions via the coupling of carboxylate ␲* to the M2 ␦ orbitals.

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

Fig. 2. The so-called IVCT or charge resonance electronic absorptions of the cations [(But CO2 )3 MM]2 (␮-O2 CCO2 )+ , where MM = Mo2 (red), MoW (purple) and W2 (blue) recorded at room temperature in THF.

With increasing length of the bridge the magnitude of the coupling falls off and as shown in Fig. 4 the spectral features of the substituted 1,4-terephthalate bridged cations have strikingly different features for M2 = Mo2 and W2 [18]. For the molybdenum complex the higher energy broad NIR resonance is typical of a Class II mixed valence ion while the sharp features of the absorption at ∼3300 cm−1 indicate that the tungsten complex is a fully delocalized Class III mixed valence ion. This again emphasizes the importance of the relative energies of the M2 ␦ orbitals with that of the bridge ␲* and the metal d-bridge ␲* overlap. For tungsten we have observed spectral features that correlate with Class III behavior to M2 to M2 distances of ∼14 A˚ [19]. The traditional treaties relating to MV ions focused on the socalled inter valence charge transfer, IVCT band in the near IR as represented in the two state model. In reality the bridge orbitals play a key role. One could not anticipate a fully delocalized mixed valence ion where the two redox centers were separated in a vacuum by 10 A˚ or more! Thus, at minimum a 3-state model is required [20]. In the case of the dicarboxylate bridge, it is the ␲* orbital of the bridge that plays the pivotal role and so these compounds all show intense MLCT absorptions in the visible NIR region of the spectrum. In solvents such as 2-methyltetrahydrofuran, 2-MeTHF, the 1 MLCT

Fig. 3. The so-called IVCT or charge resonance bands associated with the anions trans-M2 (Ti PB)2 (O2 CC6 H4 NB(C6 F5 )3 )2 − , where M2 = Mo2 (black) and W2 (red) recorded at room temperature in THF.

1579

Fig. 4. Electronic absorption spectra of the cations (bridge) [(But CO2 )3 M2 ]2 + where M2 = Mo2 (red) and W2 (blue) and the bridge is 2.5-dihydroxy-1,4-terephthalate recorded in THF solution at room temperature. The spectrum around 3000 wavenumbers is not shown because of the THF strong absorptions.

bands are temperature dependent and show vibronic features. This topic has been discussed in detail for oxalate bridged complexes [15], and a similar behavior is observed for trans-substituted complexes L-M2 -L, where L = a ␲-accepting carboxylate ligand [21]. The planar arrangement of the M2 unit with the conjugated ␲ system of the bridge/ligand is favored by M2 ␦ to bridge/ligand ␲* bonding and this is generally seen crystallographically, as shown in Fig. 5 for a thienylvinylcarboxylate compound, but thermal energy in the solution phase leads to a wide variety of rotamers and the room temperature spectra reflect this ensemble. With a lowering of temperature the molecular ensemble becomes more uniform in favoring the ground state planar geometry [21]. As the coupling of the M2 ␦ orbital and the bridge/ligand increases the (0,0) transition becomes more intense and the vibrational features are reduced as the ground and excited state surfaces start to nest. In this regard it is worth noting that this MLCT absorption for the MV ions is always sharper with less vibronic coupling – an indication that the electronic coupling is enhanced in the MV ion relative to the neutral complex [15]. This is counter intuitive on a simple MO model where introducing a positive charge would be expected to stabilize the M2 ␦ orbital leading to an increase in the energy separation of M2 ␦ and the bridge/ligand ␲* orbitals.

Fig. 5. Electronic absorption spectra of trans-W2 (Ti PB)2 (O2 CCH CHC4 H3 S)2 , where Ti PB = 2,4,6-triisopropylbenzoate, recorded in 2-methyl-THF at room temperature (red) and ·N2 temperature (blue) showing the vibronic features associated with the CO2 moiety of the thienylvinylcarboxylate ligand. The molecular structure of the complex deduced from crystallographic studies is shown in the inset.

1580

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

Fig. 6. Resonance Raman spectra of W4 OXA as the pure crystalline solid within capillaries at ca. 80 K, with 647.1 (29 mW), 568.2 (37 mW), and 514.5 (41 mW) nm excitation and spectral slit widths of 2 cm−1 . Spectra are corrected for laser power, spectrometer throughput, and the 4 scattering efficiency. The feature marked with an asterisk is due to a sudden fluctuation in laser power. W4 OXA is the complex given in the text.

Also the charge on the metal would contract the d orbitals and thereby decrease the overlap between the metal center and the bridge/ligand. However, this must be more than compensated by the resonance stabilization of the positive charge over the two redox active centers.

respectively, though in reality there is a significant mixing of all these modes. EPR spectroscopy is also valuable in tracking the location of the electron on a somewhat slower time-scale ∼ 10−9 s. A Mo2 5+ center supported by carboxylates shows g ∼ 1.9 and hyperfine coupling to 95/97 Mo, I = 5/2 of ∼27 G [25]. If the charge/hole is delocalized over both dinuclear centers the hyperfine coupling to the I = 5/2 nuclei is further reduced to ∼14 G. Similarly, a W2 5+ center with a single electron in an ␦ orbital has a central resonance at g ∼ 1.8 with a satellite spectrum due to the presence of 183 W, I = ½, 15% natural abundance Ao ∼ 52 G [25]. With delocalization over two M2 centers this hyperfine coupling is reduced to roughly half this magnitude, and a close inspection of the hyperfine spectra clearly reveals the presence of MV ions having coupling to more spin active nuclei than in the M2 4+ containing ions. This effect is nicely demonstrated in the spectra shown in Fig. 7. EPR spectroscopy is also a powerful tool in examining the polarization of a single electron bond. For example, in the MoW(O2 CBut )4 + ion the hyperfine coupling to 95/97 Mo is 44 G while that to 183 W is 30 G. Thus with respect to the homonuclear quadruply bonded ions M2 (O2 CBut )4 + , where the hyperfine coupling to Mo is 27 G and to W is 52 G, the coupling to Mo is enhanced and that to W reduced, indicative of the polarization of the single delta electron toward Mo. We can estimate Mo 4dxy ∼ 70% and W 5dxy ∼ 30% contributions to the delta bond [25]. Similarly in the mixed valence ion where two W2 units are linked by the unsymmetrical 2,6azulenedicarboxylate bridge we observed two distinct hyperfine couplings to 183 W, Ao ∼ 15 and 25 G indicative of a delocalized but polar MV ion [19].

5. Other spectroscopic techniques The molecules and MV ions under discussion show strong resonance enhancement of (MM) and the totally symmetric vibrational modes of the bridge/ligands with excitation into the MLCT absorptions. This has been well documented for oxalate [22], terephthalate [23], and 2,5-thiophenyl bridges [24]. The (MM) value of the quadruple bond shifts to lower energy upon oxidation as the bond order is reduced to form 4 to 3.5 in a valence trapped MV ion, or 3.75 in a totally delocalized MV ion. The resonance enhancement of Raman spectra of the oxalate bridged complex [W2 (O2 CBut )]2 (␮-O2 CCO2 ) is shown in Fig. 6 as the excitation wavelength moves into the MLCT transition. The vibrational modes indicated as 1 , 2 and 3 are formally the totally symmetric CO2 stretch, the CC stretch and the CO2 ␦ wagging/deformation modes,

6. Mixed valence excited states Neutral transition metal complexes can undergo photoexcitation to produce mixed valence excited states as has been well documented and studied by Zink and coworkers [26]. The neutral complexes that are the precursors to the MV ions described above are good candidates for the study of this phenomenon. Metal to bridge charge transfer places a negative charge on the bridge and a positive charge on one of the M2 centers. If the coupling of the two M2 centers is strong then the hole, the positive charge may be equally shared by both M2 centers. Similarly for compounds having a single M2 center and two ␲-acceptor ligands, the MLCT state has the positive charge on the M2 center but the negative charge could

Fig. 7. EPR spectra recorded at room temperature in THF of the cations derived from single electron oxidation of the MM(O2 CBut )4 complexes (left) and their related oxalate and perfluoroterephthalate bridged dimers of dimer cations involving Mo2 units (right).

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

1581

Fig. 8. Representation of the excited state mixed valence that arises from photoexcitation to the metal to ligand charge transfer states. The coupling in the excited state can lead to valence trapping (as depicted) or complete delocalization of the charge.

be localized on one ligand or delocalized over both. This situation is depicted in Fig. 8. For the oxalate bridged complexes [(But CO2 )3 MM]2 (␮-O2 CCO2 ) in THF the electronic coupling of the two MM centers is strong in their MV ions as shown in Fig. 2. In the neutral state the metal-tobridge charge transfer transitions occur with a minimal change in dipole moment (MM = Mo2 , MoW or W2 ) as determined by Stark spectroscopy [27]. This supports the view that both M2 centers are involved in the photoexcitation and the excited state can be viewed as a Class III, fully delocalized MV ion. However, Stark spectroscopy

for the related terephthalate bridged complex, where MM = Mo2 , has a significant dipole moment [27]. This is consistent with the weaker coupling seen in its ground state MV ion. Several of the compounds described here show dual emission from their excited states and studies of the solvent dependence of the fluorescence from the 1 MLCT states and phosphorescence from the T1 state can be informative with regard to the polar or nonpolar nature of the excited state. Rather interestingly most Mo2 containing complexes show solvent independent phosphorescence but this is not from an MLCT state but rather from the 3 MoMo␦␦*

Fig. 9. Time-resolved infrared spectra of the MLCT states of the complexes trans-M2 (Ti PB)2 (O2 CC6 H4 -4-CN)2 in the C N stretching region where M2 = Mo2 or W2 and Ti PB = 2,4,6-triisopropylbenzoate. Reproduced from Ref. [29] with permission.

1582

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

Fig. 10. Time-resolved infrared spectra of the compounds trans-M2 (O2 CMe)2 [(Pri N)2 CC CC6 H5 ]2 showing the C C stretching modes of their MLCT states. Reproduced from Ref. [29] with permission.

state. The energy of this emissive state is relative independent of the nature of the ligands and steady state emission spectra show well defined vibronic features associated with (MM) ∼ 400 cm−1 of the quadruply bonded ground state [28]. The lifetimes of the S1 MLCT states are typically in the range 1–20 ps and the T1 states range from ∼100 ns to 100 ␮s and thus both states can be monitored by time-resolved spectroscopies. The femtosecond time-resolved infrared spectra, TRIR of the compounds trans-M2 (Ti PB)2 (O2 CC4 H4 -4-CN)2 recorded in THF solution in the region of the C N stretching frequency are shown in Fig. 9 [29]. Both compounds, M2 = Mo2 and W2 show an excited state C N stretch shifted ∼65 cm−1 to lower energy relative to the ground state IR spectrum. This shift to lower energy is as predicted by density functional theory for the anions where the negative charge is delocalized over both ligands. This absorption persists for only a few ps for Mo2 consistent with the 1 MLCT lifetime. The 3 MMo␦␦* state has evidently no effect on the C N stretch. For M2 = W2 , the C N stretch in the excited state persists well into the lifetime of the T1 state. Thus we conclude that for M2 = W2 , the S1 and T1 MLCT states are delocalized. The situation is quite different for the amidinate complexes trans M2 (O2 CMe)2 [(Pri N)2 CC C-tolyl]2 which show two C C stretches in the 1 MLCT states, one shifted ∼200 cm−1 and the other ∼40 cm−1 [29]. See Fig. 10. The predicted shift for the anion is ∼120 cm−1 which leads is to propose that the negative charge is localized principally on one of the amidinate ligands with some small “spill-over” to the other. Again the TRIR spectra associated with (C C) are short lived for M2 = Mo2 consistent with the 1 MLCT state and the IR silent nature of the 3 MoMo␦␦*. For M2 = W2 the long-lived state is 3 MLCT and, as evidenced from the shift to lower energy from the two (C C) vibrations, the charge delocalization is different in the 3 MLCT state from that in the 1 MLCT state. It is more localized toward the W2 center with greater C C ␲* character.

7. Concluding remarks The MM quadruply bonded unit has proved to be very effective in the study of mixed valency for both metal centered and ligand (organic) mixed valence ions. In retrospect the systems described here are much more desirable or at least simpler than those involving d5 –d6 metal based MV ions such as in the Creutz–Taube ion [30], where low symmetry and spin orbit coupling complicate the assignment in the low energy NIR bands. This latter point has been extensively studied by Meyer [31]. The MM quadruply bonded system allows one to compare the electronic influence of 5d versus 4d metal ions in otherwise identical environments. In the study of excited state mixed valence these systems are again fortunate in yielding singlet and triplet lifetimes that are suitable for time resolved spectroscopic studies. These allow for the determination of charge localization and dynamics with time. While all the studies described briefly in this account seem rather removed from our initial objective of examining the role of MM quadruply bonded unit in organic conjugated polymers, I believe they will prove insightful with respective to charge transport, hole and electron conduction along a chain and that the time-resolved studies of the excited states will provide insight into the nature of the polymer singlet and triplet excitons. This is the matter of on-going research. Acknowledgements I thank the National Science Foundation for financial support of this work, together with the Ohio Super Computing Center for computational resources and the Ohio State Institute for Materials Research. The work of my talented coworkers and colleagues is also acknowledged for without their work nothing would have been accomplished.

M.H. Chisholm / Coordination Chemistry Reviews 257 (2013) 1576–1583

References [1] R.H. Cayton, M.H. Chisholm, J. Am. Chem. Soc. 111 (1989) 8921. [2] F.A. Cotton, C.A. Murillo, R.A. Walton, Multiple Bonds Between Metal Atoms, 3rd ed., Springer Science and Business Media, New York, 2005. [3] R.H. Cayton, M.H. Chisholm, F.D. Darrington, Angew. Chem. Int. Ed. 29 (1990) 1481. [4] D.V. Baxter, R.H. Cayton, M.H. Chisholm, J.C. Huffman, E.F. Putlina, S.L. Tagg, J.L. Weismann, J.W. Zwanziger, F.D. Darrington, J. Am. Chem. Soc. 116 (1994) 4551. [5] R.H. Cayton, M.H. Chisholm, J.C. Huffman, E.B. Lobkovsky, Angew. Chem. Int. Ed. 30 (1991) 862. [6] R.H. Cayton, M.H. Chisholm, J.C. Huffman, E.B. Lobkovsky, J. Am. Chem. Soc. 113 (1991) 8709. [7] M.H. Chisholm, A.J. Epstein, J.C. Gallucci, F. Feil, W. Pirkle, Angew. Chem. Int. Ed. 40 (2005) 6695. [8] M.H. Chisholm, C.M. Hadad, T. Spilker, J. Cluster Chem. 23 (2012) 767. [9] G.F.R. Parkin (Ed.), Electronically Coupled MM Quadruple Bonded Complexes of Molybdenum and Tungsten in Structure and Bonding, Springer-Verlag, Berlin, Heidelberg, 2010. [10] D.E. Richardson, H. Taube, Inorg. Chem. 20 (1981) 1278. [11] M. Robin, P. Day, Adv. Inorg. Chem. Radiochem. 10 (1967) 247. [12] F. Barrière, W.E. Geiger, J. Am. Chem. Soc. 128 (2006) 3980. [13] N.S. Hush, Prog. Inorg. Chem. 8 (1967) 391. [14] M.H. Chisholm, B.J. Lear, Chem. Soc. Rev. 40 (2011) 5254. [15] B.J. Lear, M.H. Chisholm, Inorg. Chem. 48 (2009) 10954. [16] B.S. Brunschwig, C. Creutz, N. Sutin, Chem. Soc. Rev. 31 (2002) 168. [17] P. Bunting, M.H. Chisholm, J.C. Gallucci, B.J. Lear, J. Am. Chem. Soc. 133 (2011) 5873.

1583

[18] M.H. Chisholm, F. Feil, C.M. Hadad, N.J. Patmore, J. Am. Chem. Soc. 127 (2005) 18150. [19] M.V. Barybin, M.H. Chisholm, N.S. Dalul, T.H. Holovics, N.J. Patmore, R.E. Robson, D.J. Zipse, J. Am. Chem. Soc. 127 (2005) 15182. [20] M.J. Ondrechen, J. Ko, L.T. Zhang, J. Am. Chem. Soc. 109 (1987) 11672. [21] B.G. Alberding, M.H. Chisholm, B.J. Lear, V. Naseri, C.R. Reed, Dalton Trans. 40 (2011) 10658. [22] B.E. Bursten, M.H. Chisholm, R.J.H. Clark, S. Firth, C.M. Hadad, A.M. Macintosh, P.J. Wilson, P.M. Woodward, J.M. Zaleski, J. Am. Chem. Soc. 124 (2002) 3050. [23] B.E. Bursten, M.H. Chisholm, R.J.H. Clark, S. Firth, C.M. Hadad, A.M. Macintosh, P.J. Wilson, P.M. Woodward, J.M. Zaleski, J. Am. Chem. Soc. 124 (2002) 12244. [24] M.J. Byrnes, M.H. Chisholm, R.J.H. Clark, J.C. Gallucci, C.M. Hadad, N.J. Patmore, Inorg. Chem. 43 (2004) 6334. [25] M.H. Chisholm, J.S. D’Acchioli, B.D. Pate, N.J. Patmore, N.S. Dalal, D. Zipse, Inorg. Chem. 44 (2005) 1061. [26] E.L. Plummer, J.I. Zink, Inorg. Chem. 45 (2006) 6556. [27] M.H. Chisholm, B.J. Lear, A. Moscatelli, L.A. Peteanu, Inorg. Chem. 49 (2010) 3706. [28] M.H. Chisholm, P.-T. Chou, Y.-H. Chou, Y. Ghosh, T.L. Gustafson, M.-L. Ho, Inorg. Chem. 47 (2008) 3415. [29] B.G. Alberding, M.H. Chisholm, J.C. Gallucci, Y. Ghosh, T.L. Gustafson, Proc. Natl. Acad. Sci. U.S.A. 108 (2011) 8152. [30] C. Creutz, H. Taube, J. Am. Chem. Soc. 91 (1969) 3988. [31] J.J. Concepion, D.M. Dattelbaum, T.J. Meyer, R.C. Rocha, Philos. Trans. R. Soc. A 399 (2008) 163. [32] M.H.Chisholm, N.J. Patmore, Dalton Trans. (2006) 3164. [33] M.H. Chisholm, R.J.H. Clark, C.M. Hadad, N.J. Hadad, Chem. Commun. (2004) 80.