Double Insertion of Thiophene Rings in Polyynediyl Chains to

Article pubs.acs.org/Organometallics

Double Insertion of Thiophene Rings in Polyynediyl Chains to Stabilize Nanoscaled Molecular Wires with [Cp*(dppe)Fe] Termini Séverine Roué,† Hiba Sahnoune,†,‡ Loïc Toupet,† Jean-François Halet,*,† and Claude Lapinte*,† †

Institut des Sciences Chimiques de Rennes, UMR 6226 CNRS-Université de Rennes 1, F-35042 Rennes, France Département de Chimie, Faculté des Sciences, Université M’Hamed Bougara, Boumerdès, Algeria

‡

S Supporting Information *

ABSTRACT: The dinuclear iron complexes [Cp*(dppe)Fe−CC− C4H2S-(CC)x-C4H2S−CC−Fe(dppe)Cp*] (2, x = 1; 3 x = 2) and [Cp*(dppe)Fe−CC−CC−C4H2S−CC−CC−Fe(dppe)Cp*] (4) were prepared in one-pot procedures from known organometallic precursors. Compound 2 was obtained from Cp*(dppe)Fe−CC− C4H2S−CC−C4H2S−CCH (6) and Cp*(dppe)FeCl (5) in 74% yield. Its relative 3, isolated in 84%, resulted from the oxidative coupling of Cp*(dppe)Fe−CC−C4H2S−CCH (7) in the presence of Cu(OAc)2 and 1,8-diazabicyclo[5.4.0]-undec-7-ene (DBU). Complex 4 was obtained from the bridging ligand 2,5-bis(trimethylsilylbutadiynyl)thiophene (8) and two equiv of 5. The new complexes were characterized by ESI-mass spectrometry, IR, multinuclear NMR, cyclic voltammetry, and Mössbauer spectroscopy. Complex 3 was also analyzed by X-ray diffraction on a single crystal. The data are consistent with a sizable metal−metal interaction across the 14- and 16-carbon atoms of the bridges. The singly and doubly oxidized forms 2(PF6)n and 3(PF6)n (n = 1, 2) were obtained by oxidation of the corresponding 18-electron iron(II) precursors with 1 and 2 equivs of ferrocenium salt, while 4 decomposed very quickly upon oxidation. The thermally stable salts 2(PF6)n and 3(PF6)n (n = 1, 2) were subjected to analyses by ESI-mass spectrometry, IR, Mössbauer, ESR, UV−vis, and NIR spectroscopies. The radical cations 2(PF6) and 3(PF6) belong to Class II of the mixed-valence Robin and Day classification with quite sizable electronic coupling parameters for large metal−metal separation (2(PF6), Hab = 262 cm−1, dFeFe = 17.7 Å; 3(PF6), Hab = 203 cm−1, dFeFe = 19.7 Å). Paramagnetic 1H NMR spectroscopy was also performed on the dicationic salts to measure the magnetic exchange between the distant spin carriers (2(PF6)2, ΔGST = −120 cm−1). The data were analyzed with the support of quantum chemistry calculations at the DFT level of theory.

■

INTRODUCTION Dinuclear complexes of the general formula [Cp*(PP)M(C C)xM(PP)Cp*] (Cp* = η5-C5Me5, PP = mono or bidentate phosphine, ((CC)x = all-carbon bridging ligand) have attracted great attention in the last two decades because of their potential applications in electronics able to serve as molecular components in the construction of electronic materials and nanotechnological devices at the molecular level.1−3 The wirelike performance of these molecules has been investigated in detail by various physical methods which have amply demonstrated that the electronic delocalization can extend over the whole organometallic unit. These molecules are recognized among the best potential candidates for incorporation into molecular junctions for two key reasons: (i) the allcarbon linkers were found to be the most efficient bridges for intramolecular charge transport,4−7 and (ii) in contrast with purely organic entities, where the highest-occupied and lowest unoccupied molecular orbitals (HOMO and LUMO, respectively) of the grafted molecules and the Fermi energy levels of the support are generally energetically separated by several eVs, organometallic termini enable facile tuning of the MO energies to facilitate junction’s conductance.8,9 © XXXX American Chemical Society

Among the Cx-bridged dinuclear systems, the homonuclear [Cp*(dppe)Fe-(CC)x-Fe(dppe)Cp*](PF6)n (x = 2−4) class exhibits an exceptionally large charge delocalization and a robustness of the oxidized species.4,7,10 Moreover, the electronic properties determined from solution-based bulk measurements have been related to electronic properties expressed at the single-molecule level.9,11 However, the development of new members with longer carbon chains in this series of molecules has not been successful, the corresponding oxidized forms generally being very unstable. Efforts to understand the reasons for the thermal instability of this family of radical cations led us to find a specific reaction of dimerization providing tetrametallic derivatives with an allcarbon assembly which differ from trendy graphite and graphene.12 Simultaneously, the instability of the radical cations with all-carbon bridges has been tried to be circumvented by insulating the unsaturated carbon linkers.13,14 Alternatively, the introduction of either aromatic rings, such as benzene,15,16 naphtalene,17 anthracene,18 and thiophene,14,19,20 or carborane Received: March 15, 2016

A

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics clusters,21 or organometallic units22 in the polyynediyl spacer has been the right alternative to prevent dimerization.12 This approach allowed the preparation of a large variety of thermally stable structures with tunable physical properties. One of our key results was for instance that the electronic communication between the remote iron centers is very similar in the all-carbon bridged complex [Cp*(dppe)Fe-(CC)4-Fe(dppe)Cp*](PF6) and its thiophene relative [Cp*(dppe)Fe−CC− C4H2S−CC−Fe(dppe)Cp*](PF6) showing that in these particular arrangements, -C4- and -(C4H2S)- fragments allow comparable charge delocalization with an excellent efficiency.3 Recent works on polyynediyl complexes of iron and ruthenium have shown that the thermal instability of the radical cation with long carbon bridges originates from the increased weight of the carbon orbitals in the description of the HOMOs, favoring the delocalization of the spin density on the carbon chain rather than on the metal ends.23 As a consequence, when the spin density on the nonsterically protected carbon atoms of the bridging ligand becomes large enough, intermolecular coupling takes place with specific formation of carbon−carbon bonds involving the γ and δ carbon atoms of two radical cations.24,25 In the quest of thermally stable molecular conductors long enough to bridge the gap between atomic-sized electrodes (≈ 2 nm), with highest conductance, we have designed a new family of molecular wires with the two main objectives: (i) to intensify the spin density on the metal termini [Cp*(dppe)Fe−CC], the Cα and Cβ atoms being sterically protected by the bulky Cp* and dppe ligands and (ii) to favor a high conductance employing in the construction of the bridge a number of carbon atoms as large as possible. With these two prior requirements, we decided to prepare some representatives of the new family of complexes [Cp*(dppe)Fe−CC−C 4 H 2 S-(CC) x C4H2S−CC−Fe(dppe)Cp*](PF6)n (x = 1, 2; n = 0−2) and to study their physical properties. We successfully synthesized and measured the synthesis and the properties of the new compounds 2(PF6)n (n = 0−2, x = 1) and 3(PF6)n (n = 0−2, x = 2, Chart 1). Moreover, to evidence the role of the thiophene rings on the stabilization of the radical cations we also prepared for the purpose of comparison the complex [Cp*(dppe)Fe−CC−CC−C4H2S−CC−C C−Fe(dppe)Cp*] (4) which contains one single thiophene ring spanning two diynediyl iron fragments. In this article, we report and discuss the main results which were obtained. Investigation of the electronic and magnetic properties of these new complexes with a wide range of spectroscopic techniques and quantum chemistry at the density-functional theory (DFT) level shows the influence of the structure of the bridges on the electronic and magnetic communication between the remote iron termini. The results are compared with the data previously obtained for the related complexes 1,19 9,16 and 10 (see Chart 1).26

Chart 1

mixture of THF/MeOH produced the dinuclear alkynyl vinylidene intermediate which was in situ deprotonated by the addition of 1 equiv of KOBut before removal of the solvents. Complex 2 was isolated in 74% yield as a pure redbrown powder. The dinuclear complex 3 was obtained from an oxidative coupling of the previously reported complex Cp*(dppe)Fe−CC−C4H2S−CCH (7).28 The iron complex 7 was reacted with Cu(OAc) 2 in the presence of 1,8diazabicyclo[5.4.0]-undec-7-ene (DBU) in pyridine.29 The reaction performed at 50 °C afforded the new complex 3 isolated in 84% yield as a pure red powder. The neutral dinuclear μ-bis(butadiynyl)-thiophene complex 4 was obtained from the bridging ligand 2,5-bis(trimethylsilylbutadiynyl)thiophene (8) and the chloro organoiron complex 5 following a three-step one-pot procedure (Scheme 1).30 The new complexes 2, 3, and 4 were isolated in an analytically pure form and were characterized by ESI-mass spectrometry, IR, multinuclear NMR, UV−vis spectroscopies and cyclic voltammetry. They are thermally stable in the solid state and in solution under argon pressure, except compound 4 which slowly decomposes in solution. As a consequence of its lack of stability, the 13C NMR spectrum of 4 could not be obtained. The 13C NMR spectra of 2 and 3 are very characteristic and unambiguously allow the determination of their structures. The chemical shift of the Cα atoms bound to the iron nuclei at δ = 151.8/156.9 for 2 and 3, respectively, is significantly larger than the value reported for 1 (δ = 139.2).19 As previously noted for bis(iron) complexes with all-carbon bridges (-CC-)x, the 13C resonances prove to be very sensitive to the extension of the carbon rich ligand spanning the two metal centers.31 Indeed, as the number of carbon atoms in the bridge increases, their electron-withdrawing effect on the

■

RESULTS AND DISCUSSION 1. Synthesis. The three neutral dinuclear complexes 2, 3, and 4 were prepared by three different routes using organometallic building blocks, which have been previously reported by our group.27,28 Compound 2 was obtained in a two-step one-pot procedure starting from the known iron complex Cp*(dppe)Fe−CC−C4H2S−CC−C4H2S−C CH (6) which contains two thiophene rings connected by an ethyndiyl fragment (Scheme 1).28 Treatment of 6 with 1 equiv of the chloro iron complex 5 in the presence of NaBPh4 in a B

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics Scheme 1. Synthesis of the Dinuclear Complexes 2, 3, and 4a

a

Key reagents: (i) NaBPh4, THF/MeOH (1/3), KOBut; (ii) Cu(OAc)2, pyridine, DBU; (iii) K2CO3, NaBPh4, THF, MeOH, KOBut.

metal increases and induces some augmentation of the π-back bonding response from the metal center, and the 13C NMR chemical upfield shift for the Cα atoms also increases. This effect is limited to the carbon atom directly bound to the metal, the vicinal Cβ atom being invariably observed between 112 and 113 ppm. The chemical shift of the sp-carbon atoms between the two thiophene rings are also sensitive to the length of the bridge. It decreases from 86.5 in 2 to 79.1/77.9 in 3, slowly getting closer and closer to the value expected for organic polyynes of the general formula R-(CC)x-R with x approaching infinity (δ ∼ 63).32 2. Molecular Structure of 3. Slow diffusion of pentane into concentrated solutions of the bis(iron) complexes 2, 3, and 4 in dichloromethane provided crystals. However, only the red crystals obtained for 3 were suitable for X-ray analysis. Diffraction parameters for this compound are available in Table S1 (Supporting Information). An ORTEP view of 3 is shown in Figure 1, while pertinent distances and angles are collected in Table 1.

Table 1. Pertinent Experimental and DFT Computed (in Brackets) Distances (Å) and Angles (deg) for 3(PF6)n (n = 0, 1, 2)

a

compd

3

3+

32+(T)a

Fe1···Fe5 Fe1-Cp*b Fe5-Cp*b Fe1−P1 Fe1−P2 Fe5−P51 Fe5−P52 Fe1−C37 C37−C38 C38−C39 Fe5−C87 C87−C88 C88−C89 C42−C43 C43−C44 C44−C94 C93−C94 C92−C93 Fe1−C37−C38 Fe5−C87−C88 C37−C38−C39 C87−C88−C89 Cp*−Fe1−Fe5-Cp*b

19.70[19.974] 1.733(2)[1.803] 1.746(2)[1.811] 2.1796(7)[2.215] 2.1990(7)[2.223] 2.1777(7)[2.217] 2.1834(7)[2.223] 1.883(2)[1.874] 1.222(3)[1.244] 1.414(3)[1.394] 1.872(2)[1.877] 1.225(3)[1.243] 1.412(3)[1.395] 1.410(3)[1.388] 1.210(4)[1.235] 1.365(4)[1.342] 1.211(4)[1.235] 1.408(3)[1.388] 178.7(2)[176.7] 176.6(2)[177.9] 175.2(3)[176.6] 174.3(3)[176.6] −3.20[−5.26]

[19.485] [1.818] [1.823] [2.239] [2.256] [2.244] [2.251] [1.848] [1.252] [1.383] [1.848] [1.253] [1.383] [1.379] [1.240] [1.335] [1.240] [1.379] [176.7] [175.5] [177.8] [177.2] [−12.75]

[19.428] [1.831] [1.839] [2.275] [2.298] [2.283] [2.295] [1.853] [1.251] [1.387] [1.856] [1.250] [1.389] [1.387] [1.235] [1.341] [1.235] [1.388] [174.5] [174.3] [177.0] [178.8] [−10.70]

Triplet state. bCentroid.

−1.0 to 1.2 V [vs standard calomel electrode (SCE)]. The voltammograms are characterized by two reversible and partially resolved waves in dichloromethane around 0 V with the ipa/ipc current ratio of unity for the two sequential electron transfers (Figure 2). A third and irreversible event is also observed at a more positive potential (ca. 1 V). The electrochemical potentials are collected in Table 2 with those reported for the closely related compound 1. As the number of carbon atoms in the bridge increases, the wave separation decreases from 4 to 2 and 3. For compounds 2 and 3, the waves almost merge into a single two-electron reversible wave. For the latter, the separation between the two waves is below the limit of resolution of the cyclic voltammogram. As a consequence, the redox potentials cannot be accurately determined as the

Figure 1. ORTEP representation of 3 at a 50% probability level. Hydrogen atoms and solvent molecules have been removed for clarity.

The two terminal iron fragments are identical within the experimental errors but are not crystallographically equivalent. They show the expected pseudo-octahedral geometry with bond lengths and angles in the previously established ranges.1,3,33 As often observed for oligothiophenes, the thiophene rings are not coplanar.34 The two heterocycles are almost planar, and the dihedral angle defined by these two planes was found to be 68.89°. 3. Cyclic Voltammetry of 2, 3, and 4 and Preparation of 2(PF6)n and 3(PF6)n (n = 1, 2). The initial scan in the cyclic voltammetry of the new complexes 2, 3, and 4 were run from C

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

[Cp2Fe](PF6) in THF at −80 °C for 6 h, and cold pentane was then added at this temperature to precipitate burgundy powders of 2(PF6)2 and 3(PF6)2 which were isolated in 85 and 83% yield, respectively. These salts are thermally stable at 20 °C under argon pressure and show the same CV as their neutral parents. Oxidation of complex 4 under the same conditions, provided a black powder which is not redox active. Attempts to identify the structure of the isolated material were unsuccessful. ESR monitoring of the reaction between 4 and the ferrocenium salts at −80 °C did not permit the recording of any signal. As a consequence of the small Kc values, treatment of 2 and 3 with a single equiv of ferrocenium salt generated in the solution the neutral and the dicationic species in addition to the targeted mixed-valence complexes. However, favored by the increase of Kc with the decrease of the temperature and possibly favorable products of solubility for these compounds, reaction of 2 and 3 with a single equiv of [Cp2Fe](PF6) allowed the isolation of the mixed-valence complexes 2(PF6) and 3(PF6) as red powders in 85 and 93% yield, respectively. These compounds can be regarded as spectroscopically pure, the ESR and IR signatures of the related neutral and dicationic partners being not observed in the spectra recorded in the solid state. 4. Electronic Structures. DFT calculations were carried out on organometallic dinuclear thiophene complexes 20/+/2+ and 30/+/2+, in order to analyze their electronic structure and bonding properties (see the Experimental Section for computational details). Results being similar overall, only those obtained for 30/+/2+ will be described and discussed here (see Supporting Information for the results obtained for 2). The structural arrangement of 3 was first optimized and compared to the available X-ray data. Computed metrical parameters match reasonably well with the available experimental values (Table 1). The largest bond length deviations are found for the FeCp* (centroid) and Fe−P distances which are computed ca. 0.07 and 0.04 Å larger than the experimental ones. As often observed for this kind of acetylide metal complex,1,3 the computed CC bond lengths are slightly overestimated by ca. 0.01 Å, with respect to the experimental ones. Computations confirm that CC bonds spanning the two thiophene rings are shorter than the outer ones tethered to the iron centers. As experimentally observed in the X-ray structure, the two iron end-caps adopt an essentially cis arrangement (Cp*(centroid)−Fe1−Fe2-Cp*(centroid) = −5.26° (the experimental value is −3.20°), the two thiophene units are not coplanar, and a torsion angle of 78.61° is computed, which compares very well with the experimental one of 68.89°. The computed iron− iron separation (19.97 Å) closely mirrors that measured experimentally (19.70 Å). Interestingly, it is important to mention that another less twisted rotamer with the two metal termini (dFe−Fe = 20.07 Å) adopting a trans arrangement and the two thiophene groups being coplanar this time (head-tohead) was also found as another gas-phase energy minimum on the potential energy surface, almost isoenergetic to 3 (0.3 kcal/ mol more stable). This might indicate that in solution or in the solid state the preferred conformer depends also on the environment (solvent, counteranion in the case of salts, or packing effects). A cis arrangement with coplanar thiophene groups (head-tohead) is computed as the most stable conformer for the monocation 3+ and dication (triplet state) 32+. In line with what is previously observed for related systems,1,3 a substantial elongation of the Fe−P bond lengths (ca. 0.05 Å) for each oxidation is also seen. A slight contraction of the Fe−C(C)

Figure 2. Cyclic voltammograms of 2 (middle), 3 (bottom), and 4 (top). Conditions are reported in Table 2.

Table 2. Electrochemical Data for Complexes 1−4a compd 1 2 3 4

E01

E02

−0.39 −0.05 −0.12 −0.04 −0.05 (ΔEp = 0.12) −0.06 0.08

E01 − E02

Kcc

0.10 0.98 0.98

0.34 0.09 0.07c

5.8 × 10 30 10

0.94

0.15

300

E03

ref 5

19 b b b

Potentials in CH2Cl2 (0.1 M [Bun4N](PF6), 25 °C, platinum electrode, and sweep rate 0.100 V s−1) are given in V vs SCE; the ferrocene−ferrocenium couple (0.46 V vs SCE)38 was used as an internal reference for the potentials measurements. bThis work. c Calculated from E02 − E01 = −(RT/F) ln Kc. a

midpoints between the anodic and cathodic peaks of the corresponding one-electron waves. However, following the method previously employed in the cases of the organoiron complexes [Cp*(dppe)Fe−CC(Si(CH 3 ) 2 ) x CC−Fe(dppe)Cp*] (x = 2−4)35 the potentials were derived from the location of the midpoint between the anodic and cathodic peaks of the two-electron wave and the distance between them (ΔEp), provided that kinetics of the electron-transfer processes does not affect the CV response.36 The observation of negligible variations of ΔEp as a function of the scan rate below 1 V s−1 indicates that this condition is fulfilled and that fast electron transfer processes takes place. Thus, values of 0.09 and 0.07 V were obtained for the differences ΔE0 = E02 − E01 for complexes 2 and 3, respectively (for details, see ref 35 and Figure S1 therein). As reported by Launay and co-workers, Kc can also be determined by a titration method.37 This method was applied to samples of 3 using various equivalents (0 to 2) of [(C5H5)2Fe](PF6). The results (10 < Kc < 15) confirmed the data obtained from the CV measurements. The molar fractions of the [Fe(II)−Fe(II)], [Fe(II)−Fe(III)]+, and [Fe(III)−Fe(III)]2+ species in solution calculated using the value of Kc reported in Table 2 are 0.13/ 0.74/0.13 and 0.19/0.62/0.19 for the MV complexes 2(PF6) and 3(PF6), respectively. In comparison with complex 1, the first oxidation potential of the new complexes 2, 3, and 4 are shifted toward less negative values by 0.27, 0.34, and 0.33 V, respectively.19 These data may reflect the electron-withdrawing effect of the -CC- alkynyl fragments and also the weakening of the metal−metal interactions. It is noteworthy that the potentials of these compounds are all far below the potential of the redox couple [Cp2Fe/Cp2Fe+]. Therefore, their one-electron and two-electron oxidations can be performed using ferrocenium salt. The neutral dinuclear complexes 2 and 3 were reacted with two equivs of D

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics (ca. 0.03 Å) is noted for the first oxidation. No change is observed upon the second one. The computed Fe-Cp*(centroid) and outer CC bonds adjacent to the metal centers slightly lengthen in the monocation by ca. 0.01 Å on average but do not change in the dication. The CC bonds framed between the thiophene groups hardly change upon oxidation. Other conformers were envisaged and optimized. Unfortunately, no other energy minima could be trapped within our “gas-phase” calculation conditions. This does not mean that there are none, taking into account the very low energy barrier for rotation around a triple bond.39 Structural change upon oxidation can partly be understood by examining the nodal properties of the HOMOs of the neutral system 3 (see Figures 3 and 4). As usually observed for

oxidized than the former (see section 3). The second IPs are similar: 11.27 and 11.25 eV for 2+ and 3+, respectively. 5. IR Spectroscopy. The IR spectra were recorded for samples of 2(PF6)n, 3(PF6)n (n = 0−2, Figure 5), and 4 on

Figure 5. IR spectra in Nujol (KBr plates) of 2(PF6)n (top) and 3(PF6)n (bottom, n = 0, 1, 2). Figure 3. DFT molecular orbital diagram of 3. Fe (left)/carbon chain (middle)/sulfur(right) percentage contributions are given in italics.

powders (Nujol mull) and in solution of dichloromethane. The frequencies of the νCC bands are collected in Table 3. The Table 3. Experimental IR νCC Bands for 1(PF6)n, 2(PF6)n, 3(PF6)n, and 4 (n = 0, 1, 2) in Nujol and (CH2Cl2) compd

n=0

n=1

1(PF6)n

2054, 2039 (2041) 2176, 2032 (2150, 2026) 2176, 2125, 2031 (2124, 2021) 2131, 1991 (2125, 1987)

1977, 1910 (1983, 1927) 2153, 2017, 1931 (2012, 1929) 2162, 2124, 2019, 1948 (2012, 1939)

2(PF6)n 3(PF6)n 4

Figure 4. Plots of the HOMO (−3.59 eV), HOMO−1 (−3.71 eV), LUMO (−1.94 eV), and LUMO+1 (−1.92 eV) are shown. Contour values are ±0.03 (e/bohr3)1/2.

a

n=2

ref

1941 (1950)

19

2176, 1943 (1945)

a

2178, 2129, 1953 (2180, 2132, 1949)

a a

This work.

valence modes of the CC triple bonds directly attached to the iron nuclei are more intense and appear at lower frequencies than the CC triple bonds in the vicinity of the thiophene rings. These latter bands are so weak, especially in the case of the homovalent complexes (n = 0 and 2), that they are not always visible in the spectra (see, for example, the spectrum of 2 and 2(PF6)2 in Figure 5). When visible, the νCC bands corresponding to the CC triple bond between the two thiophene rings of the complexes 2 and 2(PF6)2, are not sensitive to the oxidation state of the metal (ν = 2176 cm−1 in both cases in nujol), and for symmetry reasons can be assigned to the antisymmetric mode of vibration. In contrast, a band much more intense can be observed at a lower frequency (ν = 2153 cm−1) in the heterovalent compound 2(PF6). Probably, this band corresponds to a

this kind of metal-alkynyl systems,1−3 the HOMOs of complex 3 substantially separated from the LUMO by 1.65 eV are part of the “t2g” set expected for pseudo-octahedral metal centers, π in character, heavily localized on the Fe centers and on the adjacent ethynyl groups, and here, to a lesser extent, on the thiophene rings. They are, in turn, mainly Fe1−C37 and Fe5− C87 antibonding and C37−C38, and C87−C88 bonding in character. Consequently, the oxidation of 3 leads to some shortening of the Fe−C distances and a slight lengthening of the adjacent CC bonds (vide supra). The first adiabatic ionization potential (IP) for the cation 3+ is computed to be 4.74 eV. For comparison, that calculated for 2+ is 4.43 eV. This in agreement with the electrochemistry measurements, which show that the latter is more easily E

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics symmetric mode of vibration of the cation suggesting that the symmetry of the molecule is broken in the mixed-valence state. Concerning the three compounds of the 3(PF6)n family, two vibration modes are clearly observed for the butadiendiyl fragment linking the two thiophene rings. The frequencies weakly decrease from the neutral to the monocationic complexes and increase to a similar extent from the monocationic to the dicationic derivatives, suggesting that a small reorganization of the bonding of the spacer is associated with the oxidation states of the metal centers. From these data, it seems that the electronic delocalization is more favored along the carbon bridge in the mixed-valence state than in the homovalent states. The spectra run in CH2Cl2 solution are very similar. The vibration modes corresponding to the CC triple bonds which link the thiophene rings are of weak intensity and sometimes difficult to observe, but the Fe−CC vibrations clearly show the same multiplicity as that of the spectra collected for samples in the solid state. However, one can note significant shifts toward the lower energies (ca. 5−9 cm−1) suggesting that free rotation in solution probably favors the electronic delocalization along the bridge. The IR spectra of the neutral complexes 2 and 3 exhibit also one strong band (2032 and 2031 cm−1, respectively) characteristic of the antisymmetric mode of vibration of the Fe(II)-CC triple bonds.40 Similarly, the spectra of the dicationic compounds 2(PF6)2 and 3(PF6)2 display a band (1943 and 1953 cm−1, respectively) typical of the same mode of vibration of the Fe(III)-CC triple bonds.40−42 These results, indeed, seem to be supported by the X-ray structures of 3 and the computed geometries of 3n+ (n = 0−2) which show CC bond distances in the C4 unit spanning the two thiophene rings shorter (and almost constant upon oxidation) than those of the CC triple bonds tethered to the metal atoms (see Table 1). Additionally, two frequencies, one of each being very intense, are computed at 2035 cm−1 (5476 km/mol) and 2040 cm−1 (109 km/mol) for 3. They are associated with the Fe(II)-CC vibrators. Two weak frequencies are also calculated at 2118 cm−1 (94 km/mol) and 2162 cm−1 (4 km/mol), corresponding to the vibrations of the CC bonds framed between the two thiophene rings. In contrast to the spectra of the neutral complexes, the spectra of the mixed-valence complexes 2(PF6) and 3(PF6) exhibit two bands corresponding to the Fe−CC vibrators (2017/1931 and 2019/1948 cm−1, respectively). In the range 2100−1800 cm−1, the IR spectra of the monocations 2(PF6) and 3(PF6) can be roughly regarded as the sum of the spectra of the corresponding homovalent Fe(II)−Fe(II) and Fe(III)− Fe(III) complexes. These data clearly establish that the symmetry of the complexes is broken for these two monocations and that the rate of the intramolecular electron transfer is slow on the fast IR time scale (10−12 s). However, the frequencies of the Fe−CC bands are slightly shifted toward the low energies with respect to their homovalent relatives suggesting a diminution of the CC bond order in the mixedvalence state. In other words, the canonical form B depicted in Scheme 2 might significantly contribute to the description of the electronic structures of 2(PF6) and 3(PF6). The IR spectrum of the new iron diynyl complex 4 shows two bands at 2131 and 1991 cm−1 in the solid state. The lowest energy band corresponds to a νCC vibration mode of the alkynyl-iron fragment, while the band on the high energy side

Scheme 2. Selected Possible Mesomeric Structures for 2(PF6) and 3(PF6)

corresponds to the vibration mode of the alkynyl-thiophene moiety. 6. Mö ssbauer Spectroscopy. The Mössbauer spectra of the complexes 2(PF6)n and 3(PF6)n (n = 0−2) were run at 80 K and least-squares fitted with Lorentzian line shapes (Figure 6

Figure 6. 57Fe Mössbauer spectrum of 3(PF6) at 80 K.

and Figure S2 in the Supporting Information).43 The isomeric shift (IS, δ) and quadrupole splitting (QS) parameters are given in Table 4 with the data previously reported for the reference Table 4. Mössbauer Parameters for the Complexes 1(PF6)n, 2(PF6)n, and 3(PF6)n (n = 0−2) QS (IS) mm s−1 vs Fe, 80 K compd T (K) 1 1(PF6) 1(PF6)2 2 2(PF6) 2(PF6)2 3 3(PF6) 3(PF6)2

Fe(II)

Fe (averaged)

Fe(III)

1.984 (0.255) 1.444 (0.212) 1.024 (0.181) 1.970 (0.241) 1.934 (0.252) 1.983 (0.246) 1.961 (0.257)

0.901 (0.241) 0.922 (0.231) 0.905 (0.233) 0.914 (0.235)

relative areas (%) 100 100 100 100 50/50 100 100 50/50 100

compounds 1(PF6)n (n = 0−2). The QS and IS parameters obtained for the homovalent derivatives are close to the data obtained for many mononuclear Fe(II)-CC and Fe(III)-C C complexes in the Cp*(dppe)Fe series.1,3 The QS parameters found for the dications 2(PF6)2 and 3(PF6)2 are typical of low spin d5 Fe(III) centers.3,44 These compounds contain two unpaired electrons, and it can be anticipated that the samples may contain a mixture of molecules in the singlet state (S = 0) and in the triplet state (S = 1). This is in accord with the theoretical calculations that show similar energy for the antiferromagnetic and ferromagnetic states (vide supra). In a previous work, it has been shown that a smaller value of QS is expected for paramagnetic compounds or pure triplet states (ca. 0.8 mm/s), while a larger value is characteristic of a singlet state (ca. 1.1 mm/s).45 F

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics Therefore, the QS values found for 2(PF6)2 (QS = 0.922 mm/ s) and 3(PF 6 ) 2 (QS = 0.914) suggest that a small antiferromagnetic exchange should take place between the remote spin carriers in these two compounds with an energy gap between the singlet and triplet states larger for 2(PF6)2 than 3(PF6)2. The 57Fe Mössbauer spectra of the heterovalent complexes 2(PF6) and 3(PF6) display two distinct doublets with relative spectral absorption areas in the 1:1 ratio, diagnostic of localized valence. It is important to report that, within the accuracy of the fits of the experimental spectra, an increase of the line broadening cannot be detected for the MV complexes, with respect to those of the homovalent parent derivatives. Consequently, the exchange rate constant ke < 106 s−1 for the intramolecular electron transfer in 2(PF6) and 3(PF6).46 This behavior strongly contrasts with the previous observation for the MV complex 1(PF6). Indeed, the spectrum of 1(PF6) exhibits a single doublet with a QS parameter close to the averaged values found for the homovalent Fe(II)−Fe(II) and Fe(III)−Fe(III) derivatives as shown in Table 4 indicating that ke > 109 s−1.4,10 In previous studies on MV complexes containing the Cp*(dppe)Fe termini, it has been found that some solid samples can contain both valence trapped and valence detrapped molecules depending on the conformational arrangement of the radical cations.15,19,47 If one considers that the spectra were recorded using powdered samples obtained by precipitation of the salt in solution at low temperature, it is very likely that various conformers of the radical cations are present in the powder. Therefore, we tend to think that the charge is localized on the Mössbauer time scale for all the represented conformers of 2(PF6) and 3(PF6). 7. Glass ESR Spectroscopy. The X-band ESR spectra of the radical cations 2(PF6) and 3(PF6) were recorded at 66 K in a rigid glass (CH2Cl2/C2H4Cl2, 1:1). The spectra of the monocations display three features characteristic of one lowspin d5 Fe(III) ion in a pseudo-octahedral environment (Figure 7).3 The values of the g-tensor components were extracted from the spectra and collected in Table 5. Note that the extra and broad feature observed in the spectrum of 3(PF6) (marked with an asterisk) belongs to the dication 3(PF6)2 also present in solution. The ESR parameters are consistent with a small degree of delocalization of the odd electron in the MV 2(PF6) and 3(PF6) at 66 K. Indeed, for MV compounds of a homogeneous series, the anisotropy of the g-tensor (Δg = g1 − g3) decreases as the rate of the intramolecular electron transfer increases.48 This property has been verified without exception in several cases in the Cp*(dppe)Fe series.1,3,19,47 The Δg value is less than 0.01% smaller in 2(PF6) than in 3(PF6) but 220% larger in 2(PF6) than in 1(PF6). This is fully consistent with the fact that 1(PF6) is a fully delocalized Class-III MV complex, while 2(PF6) and 3(PF6) belong to the localized Class-II MV derivatives. These data also evidence that the presence of an extra CC triple bond in 3(PF6) with respect to 2(PF6) has a sizable but very small effect on the rate of the intramolecular electron transfer (ET). It is also interesting to emphasize that the anisotropy (Δg) is 27% and 14% smaller in 3(PF6) than in 9(PF6) and 10(PF6), respectively,16,26 suggesting that the intramolecular ET is faster through 17 bonds in 3(PF6) than through 13 bonds in 9(PF6) and 10(PF6). The determination of the electronic coupling parameters (Hab) by analysis of the intervalence charge transfer (IVCT) band (see section 10)

Figure 7. X-band ESR spectra of 2(PF6) (top) and 3(PF6) (bottom) at 66 K in a 1:1 mixture of dichloromethane and 1,2-dichloroethane.

Table 5. ESR Parameters for 2(PF6)n, 3(PF6)n (n = 1, 2), and Some Related MV Complexes at 66 K compd 1(PF6) 2(PF6) 2(PF6)2 3(PF6) 3(PF6)2 7(PF6) 9(PF6) 10(PF6)

g1

g2

2.113 2.039 2.330 2.036 2.13 (ΔHpp = 250 2.334 2.034 2.12 (ΔHpp = 333 2.366 2.036 2.418 2.134 2.380 2.035

g3 2.006 1.989 G) 1.986 G) 1.986 1.976 1.983

giso

Δg

2.053 2.118

0.107 0.345

2.128

0.348

2.129 2.176 2.133

0.380 0.442 0.397

ref 19 this this this this this 16 26

work work work work work

supports this conclusion for 9(PF6), but not for 10(PF6). This experimental discrepancy might originate from the somewhat different nodal properties of the partially occupied MOs between thiophene- vs phenyl-containing compounds. As previously reported, the giso and Δg values are significantly smaller for complexes containing alkynyl-thiophene fragment like 7(PF6) (giso = 2.129, Δg = 0.380) than for those bearing alkynyl-aryl substituents (i.e., [Cp*(dppe)FeCC−C6H5](PF6), giso = 2.157, Δg = 0.489).41,49 For this reason, compounds 2(PF 6 ) and 3(PF 6 ) do not constitute a homogeneous series with 9(PF6) and 10(PF6). Nevertheless, the ESR tensors also indicate that with respect to a phenyl substituent, a thiophene ring favors the mixing of the d orbitals of the metal with the sp and sp2 carbon orbitals of the bridging ligand. Contrary to experiments, spin density analysis of the unpaired electron distribution in complex 3+ shows a significant delocalization all over the metal−carbon backbone, with Mulliken atomic spin densities of 0.30 e on each iron center and 0.40 e on the carbon bridge (see Figure 8). As said earlier, we were not able to trap any electron-localized species on the computed potential energy curve of 3+. In agreement with the electron delocalized Class-III picture of the computed complex 3+, the calculated components of its g tensor are relatively close G

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

below and NMR active in the temperature range where the usual solvents are liquid.27 1H NMR spectra were run for 2(PF6)2 and 3(PF6)2 in acetone d6 between 30 and −12 ppm. The spectra of the dications recorded at 20 °C are well resolved, and as expected for paramagnetic compounds, the resonances are shifted with respect to those of the neutral complexes. The most notable features are the shifts of the Cp* at δ −8.26 and −9.43 for 2(PF6)2 and 3(PF6)2, respectively, and the Hb protons of the thiophene rings at δ 20.39 and 31.49, respectively (Ha and Hb protons of the thiophene ring are attached to the carbon atom located on the side of the alkynyliron or on the opposite side, respectively). Protons Ha are probably out of the explored window, below −12 ppm. The two nonequivalent protons of the −CH2− groups of the dppe are found at δ −2.30 and 5.9 for 2(PF6)2 and at −3.21 and 5.9 for 3(PF6)2. For paramagnetic compounds, the observed isotropic shift may arise from contact and/or dipolar interactions (eq 1).50 In the case of simple mononuclear systems with one unpaired electron (S = 1/2), both contact and dipolar terms are expected to have inverse temperature dependence, but for compounds with S > 1/2, the zero-field splitting can also lead to a dipolar shift with a T−2 temperature dependence.51 The presence of simultaneous contact and dipolar contributions to the observed isotropic shifts can be established by the detection of curvature in a Curie plot.52 Observed isotropic shifts plotted against 1/T for compounds containing two (S = 1) and three (S = 3/2) [Cp*(dppe)Fe]+ units show in all cases linear relationships, indicating that the Curie law is obeyed for this family of compounds and that the isotropic shift is essentially contact in origin.53 However, for compounds in which the singlet/triplet energy gap is large enough, population of the ground state can significantly increase as the temperature decreases, thereby inducing a deviation from the Curie law. The variation of the magnetic susceptibility, χ, is then given by eq 2, which is derived from the van Vleck equation:50,54

Figure 8. Spatial distribution of the spin density computed for 3+ (isocontour values: ± 0.002 e/bohr3). Numerical parameters indicate Mulliken atomic spin density.

to 2 (Table S2, Supporting Information) with an anisotropy of 0.108 only, far from the experimentally measured value of 0.379 (see Table 5). It is noteworthy that the spin density distributed over the inner C4 chain framed between the thiophene rings is rather weak (≤0.03 e). Indeed, this is the reason why such species do not dimerize, contrary to related systems where spin density on carbon atoms is important.12,25 A similar spin density is computed for the inner C2 chain in 2+ (Figure S4, Supporting Information). When recorded in glass and in the same conditions as those of the MV complexes, the ESR spectra of the corresponding dications 2(PF6)2 and 3(PF6)2 are flat. Similar observations were also reported for 1(PF6)2 and for other rare examples of bis-iron(III) complexes in the Cp*(dppe)Fe series.15,49 This behavior is sometimes the consequence of a strong antiferromagnetic coupling in these compounds. However, measurements on powdered samples of 2(PF6)2 and 3(PF6)2 show weak signals (Figure 9). It is also mentioned that the

δobs = δiso + δdia = δcontact + δdipolar + δdia

(1)

χ = C /T[3 + exp( −2J /kT )]

(2)

1

Variable-temperature H NMR spectra were run for 2(PF6)2 and 3(PF6)2 between 193 and 293 K. As previously observed for related examples, the presence of the unpaired electrons predominantly affects the chemical shifts of the methyl groups of the Cp* ligands for which the line broadening at half height increases from 96/116 Hz at 293 K to 564/741 Hz at 193 K for 2(PF6)2/3(PF6)2.4,6,27,55 Plots of the experimental chemical shifts (δiso = δobs - δdia) against 1/T display a deviation from linearity in the case of 2(PF6)2, while the Curie law is perfectly obeyed (r2 = 0.999) in the case of 3(PF6)2 in the explored range of temperatures (Figure 10). This linear dependence on 1/T indicates a fairly constant concentration of the paramagnetic species over the investigated temperature range. In the case of 2(PF6)2, the temperature dependence of the C5Me5 protons could be fitted with eq 2. The best fit was obtained for J = −60 ± 2 cm−1. The negative sign for the J value indicates that the low spin state (S = 0) is the ground state and that the unpaired electrons are antiferromagnetically coupled. The triplet excited state lies above it with an energy separation equal to ΔGST = 2J = −120 cm−1. Considering the distance between the Fe(III) spin carriers (ca. 17.7 Å; a value of 17.47 Å is DFT computed for 22+), the intramolecular magnetic exchange is remarkably large. Comparison can be made with

Figure 9. X-band ESR spectra of powdered samples of 2(PF6)2 and 3(PF6)2 at 66 K.

forbidden transition Δms = ±2 characteristic of the triplet state could not be observed around 1500 G due to the weakness of the signal. However, this transition was observed in the spectrum of 1(PF6)2 for samples in the solid state.49 To complement the ESR measurements, magnetization experiments were performed on microcrystalline samples of 2(PF6)2 and 3(PF6)2 between 4 and 300 K. The thermal variation of the χMT product with temperature could not be fitted properly with the Bleaney−Bowers law. Apparently, the samples contained substantial traces of ferromagnetic impurities precluding any modeling of the experimental data. To overcome this difficulty, the magnetic properties of the two biradicals 2(PF6)2 and 3(PF6)2 were further investigated by 1H NMR in solution. 8. Paramagnetic 1H NMR Spectroscopy. It is one of the particularities of the iron(III) complexes of the [Cp*(dppe)Fe]+ series to be ESR active at liquid nitrogen temperature and H

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

characteristic data are collected in Table 6. For all compounds, broad absorptions bands were observed in the UV range (220− Table 6. Absorptions for Complexes 2(PF6)n, 3(PF6)n, and 4 (n = 0−2) in CH2Cl2 at 25 °C compd 2 2(PF6) 2(PF6)2 3 3(PF6)

Figure 10. Plots of δiso of the protons of the Cp* ligands as a function of 1/T (K) for 2(PF6)2 and 3(PF6)2 (experimental data, squares; calculated values, colored lines; the black lines show the deviation from the linear Curie law).

3(PF6)2 4

the closely related complexes [Cp*(dppe)Fe−CC-1,4C5H2S−CC−Fe(dppe)Cp*](PF6)2 (1), [Cp*(dppe)Fe− CC-1,4-C6H4−CC−Fe(dppe)Cp*](PF6)2, and [Cp*(dppe)Fe−CC-(1,4-C6H4)2-CC−Fe(dppe)Cp*](PF6)2 (9) for which the metal−metal distances were estimated to be 11.6, 11.6, and 16.1 Å and antiferromagnetic coupling were found with energy gaps between the spin states of −300 cm−1, −380 cm−1, and −2 cm−1, respectively.49,55,56 The energies of the triplet state (high-spin (HS) ferromagnetic state), and the broken symmetry singlet (BS antiferromagnetic state) of 22+ and 32+ were computed and compared. It turns out that using the expression J = EBS − 3 EHS,57 comparable exchange magnetic couplings of −274 cm−1 and −242 cm−1 are computed at the B3LYP level of theory, respectively, favoring the antiferromagnetic state for both species. In agreement with Mössbauer measurements (see above), the singlet−triplet energy gap is slightly larger for the former. Spin densities computed for the antiferromagnetic state (top) and ferromagnetic state of 32+ are pictured in Figure 11. In both cases, unpaired spin density resides principally at the metal centers (somewhat higher in the triplet state), but delocalization onto the bridge is still significant. 9. UV−Vis Spectroscopy. The UV−vis spectra for complexes 2(PF6)n, 3(PF6)n, and 4 (n = 0−2) were recorded between 220 and 800 nm at 20 °C in CH2Cl2. The

absorption λ (nm) (ε × 10−3 (M−1 cm−1)) 230 (73.6), 328 (18.8), 228 (57.0), 258 (32.3), 702 (5.4), 870 (5.6) 232 (62.9), 372 (26.6), 251 (56.0), 325 (26.6), 238 (54.5), 282 (57.0), 724 (9.2), 858 (8.2) 237 (63.3), 265 (49.7), 846 (9.3) 242 (47.4), 274 (37.3),

496 (32.4) 368 (15.3), 512 (7.4), 556 (8.4), 504 (14.2), 576 (11.0), 856 (19.8) 532 (50.1) 354 (27.9), 534 (20.3), 568 (19.2), 388 (24.3), 504 (14.2), 576 (11.0), 332 (32.7)

400 nm) with two or three maxima of shoulders assigned to π → π* transitions of the ligands. The spectra of the neutral compounds 2 and 3 show also absorptions at 496 and 532 nm, respectively, which correspond to a metal-to-ligand chargetransfer (MLCT) bands.58 Time-dependent density functional theory (TD-DFT) calculations on 3 show a rather intense adsorption band at 539 nm involving HOMO−LUMO region transitions, i.e., MLCT. These bands are responsible for the orange color of the neutral complexes but also exist in the spectra of the deeply colored oxidized forms. In the mono- and dioxidized species, the MLCT bands are hypsochromically shifted. The spectra of the doubly oxidized compounds display additional bands at 856 and 846 nm for 2(PF6)2 and 3(PF6)2, respectively, assigned to LMCT transitions.42 The spectra of the MV complexes can be regarded as the addition of the spectra of the Fe(II)−Fe(II) and Fe(III)− Fe(III) complexes. However, a red shift of the MLCT transitions with respect to the neutral parent, and a blue shift of the LMCT transitions with respect to the doubly oxidized compounds can be noted. This is consistent with the Class-II character of these MV complexes which present a sizable degree of electronic coupling.47 10. NIR Spectroscopy. The NIR spectra of 2(PF6)n and 3(PF6)n (n = 0−2) dissolved in CH2Cl2 were recorded at 25 °C (Figure 12). As expected for Fe(II) derivatives, the spectra of the neutral complexes do not contain any absorption bands in the NIR range.47 In contrast, the spectra of the dicationic complexes 2(PF6)2 and 3(PF6)2 show a band of weak intensity [ν = 5640 cm−1 (ε = 315 M−1 cm−1) and ν = 5450 cm−1 (ε = 325 M−1 cm−1), respectively] corresponding to the forbidden ligand field (LF) transition from the SOMO-2 to the SOMO invariably observed for the mononuclear Fe(III) and dinuclear Fe(III)−Fe(III) complexes in the Cp*(dppe)Fe series.42 Finally, the spectra of the MV complexes 2(PF6) and 3(PF6) display broad absorption bands of moderate intensity (ε = 3500 M−1 cm−1) with complex shapes (Figure 11). Generally, the deconvolution of experimental NIR spectra found for dinuclear MV complexes of the Cp*(dppe)Fe series is accurately achieved using three Gaussian functions.1,3,47 In the case of the MV complexes 2(PF6) and 3(PF6), a good fit of the experimental spectra was obtained using two Gaussian components only (bands A and B) and the tail of the strong absorptions in the visible range (D). It cannot be excluded that a third component with a maximum around 10000 cm−1 is hidden behind band D. An extra component (band C) has been added to the fit to represent the calculated contribution of the

Figure 11. Spatial distribution of the spin density computed for the antiferromagnetic state (top) and ferromagnetic state of 3 2+ (isocontour values: ± 0.002 e/bohr3). Numerical parameters indicate Mulliken atomic spin density. I

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics

homovalent species 2(PF6)2 and 3(PF6)2 also present in the solution (with the NIR transparent species 2 and 3) due to the weak comproportionation Kc constant. It is clear that this component has no significant effect on the quality of the fit (Table 7). Table 7. NIR Data for 1(PF6), 2(PF6), and 3(PF6) in CH2Cl2 at 25 °C band

1(PF6)

A: M→M B: M→M A: M→M B: M→M C: LF A: M→M B: M→M C: LF

2(PF6)

3(PF6)

ν (cm−1)

ε (M−1 cm−1)

Δν1/2obs (cm−1)

Δν1/2calc (cm−1)

5030 6450 4800 7170 5640 5000 6980

25000 11700 3100 1870 100 2210 2230 5450

1500 3200 3400 4100 1860 3400 4040 100

3400 3900 3340 4069 3398 4015 1860

The presence of bands A and B in the NIR range, which do not exist in the spectra of the corresponding homovalent species, clearly confirms that these new complexes constitute original examples of weakly coupled Class-II organometallic MV compounds, according to the Robin and Day classification.59 For weakly coupled systems of the [Cp*(dppe)Fe−C C-B-CC−Fe(dppe)Cp*][PF6] (B = carbon-rich bridge), the spectroscopic parameters for the intervalence charge transfer bands obey the two-level Hush model.60 According to this model and within the weak interaction limit, the photoinduced electron transfer gives rise to absorption bands with Gaussianlike shape, and the full width of the band at half height obeys eq 3: (Δν1/2)calc = (2310 × νmax )1/2

(4)

E IVCT(2) ≈ λ + E LF(1)

(5)

E IVCT(3) ≈ λ + E LF(2)

(6)

Hab = (2.06x10−2 /dab)(εmax νmax Δν1/2)1/2

(7)

Moreover, Meyer and co-workers have shown that inorganic MV complexes with d6/d5 metal centers in pseudo-octahedral environments are characterized by three transitions.61 The energies of the three optical metal−metal electron transfers (IVCT) are roughly related to the energies of the two LF transitions, according to eqs 4−6. They have the same halfwidth, and their intensity decreases rapidly as their energy increases.61,62 The energy of the LF(1) transition which corresponds to the energy difference between the SOMO and SOMO-2 is close to 2000 cm−1.42 This value is not experimentally accessible but can be estimated by DFT calculations.42 In a limited number of examples, the LF(2) energy can be spectroscopically measured when the IVCT bands are not too intense.16 In the case of the MV complexes 2(PF6)2 and 3(PF6)2, one can assume that probably a third band of small intensity corresponding to the transition with EIVCT(3) energy is hidden below the tail of the visible absorption (D). In the two-state model approximation, the NIR spectra of the Class-II MV complexes are expected to be solvent sensitive. This feature has been observed for all of the models studied in our group containing the Cp*(dppe)Fe unit with only two exceptions where the mediating bridges B are 1,3-phenyl and 3,5-pyridine. In these MV compounds, we discovered that the LMCT transition is less energetic than the MMCT transition. As a consequence, the direct metal−metal coupling does not occur (Hab = 0), and only the coupling through the bridge and the reactant and product diabatic states is effective.46 In contrast with the related MV complexes which obey the twolevel model, these two compounds present NIR transitions almost independent of the solvent polarity. For this reason, the influence of the solvent polarity was carefully investigated for the MV complexes 2(PF6) and 3(PF6). The NIR spectra of the MV complexes 2(PF6) and 3(PF6) were also recorded in 1,2-dichlorobenzene, tetrahydrofuran, acetone, and methanol. Comparison of the spectra obtained in these various solvents clearly establishes a solvatochromic behavior for the envelope of the broad NIR absorption. Spectral data extracted from the fits of these spectra reveal that both Gaussian bands A and B are shifted toward higher energies when the polarity of the solvent defined by 1/εop − 1/ εs (εop and εs are the optical and static dielectric constants of the solvent, respectively) increases (Table 8).63,64 A plot of the

Figure 12. NIR spectra of 2(PF6) (top) and 3(PF6) bottom in CH2Cl2 and proposed deconvolution.

compd

E IVCT(1) ≈ λ

Table 8. Solvatochromy of the Intervalence Bands for 2(PF6) and 3(PF6) at 25 °C 2(PF6) νmax (cm−1)

(3) J

3(PF6) νmax (cm−1)

solvent

1/εop−1/εs

band A

band B

band A

band B

1,2-dichloro-benzene THF CH2Cl2 acetone methanol

0.315 0.373 0.383 0.493 0.536

4600 4750 4800 5200 5350

6700 7120 7170 7650 7850

4700 4960 5000 5550 5600

6650 6950 6980 7800 8200

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics energy of the NIR transitions as a function of (1/εop − 1/εs) gives straight lines (r2 = 0.99) with intercepts of 3565 and 4600 cm−1 for the bands A and B, respectively, of the two compounds 2(PF6) and 3(PF6) (Figure 13). The intercept is

Hab parameters was achieved using the full experimental NIR bandwidth, which encompasses all-sub-bands. Recalculation of the electronic coupling as done by Launay et al. provides data very similar to those found by these authors (Hab = 532 cm−1 for 2(PF6) and 468 cm−1 for 3(PF6)) indicating than these compounds are at the highest limit of molecular conductance known until now of electronic coupling at the 2 nm scale.

■

CONCLUSIONS In this contribution, the syntheses in good yields of the new nanoscaled molecular wires complexes [Cp*(dppe)Fe−C C−C4H2S-(CC)x-C4H2S−CC−Fe(dppe)Cp*](PF6)n (x = 1 (2), 2 (3); n = 0−2) are described. These complexes were characterized by various means including IR, multinuclear NMR, cyclic voltammetry, 57Fe Mössbauer, ESR in rigid glass, UV−vis, and NIR spectroscopy. In addition, an XRD analysis was also carried out for the neutral complex [Cp*(dppe)Fe− CC−C4H2S-(CC)2-C4H2S−CC−Fe(dppe)Cp*]. The experimental data were analyzed with the support of DFT calculations on the full structures. This study highlights that despite an extended π-system from one metal end to the other over 14 or 16 carbon atoms, the oxidized species 2(PF6)n and 3(PF6)n (n = 1, 2) which contain one or two unpaired electrons are thermally very stable in the solid state and also in solution. This remarkable behavior is explained by the significant but rather weak carbon spin density (≤0.03 e) over the inner C4 chain framed between the thiophene rings. Indeed, the presence of a larger spin density on the carbon atoms of the bridge without steric protection would favor dimerization of the species with an open valence shell as previously observed. Importantly, the spectroscopic data show the localization of the charge in the MV states even at the slow time scale of the Mössbauer spectroscopy. These MV complexes belong to the Class II of MV derivatives as defined by Robin and Day, and the Hush model is rather well obeyed. Considering throughspace distance between the redox centers, the experimental Hab parameters are large indicating than these compounds are at the highest limit of molecular conductance at such a distance. In the case of the dicationic complex 2(PF6)2, the deviation from the Curie law of the variation of the magnetic susceptibility over 1/T indicates that the unpaired electron centered on the remote iron(III) cations is antiferromagnetically coupled with an energy separation of −120 cm−1 between the singlet ground state and the triplet excited state.

Figure 13. Solvent dependence of the NIR bands A (bottom) and B (top) for 2(PF6) (dotted line, empty square) and 3(PF6) (solid line, full square).

interpreted as the Franck−Condon inner sphere optical activation energy (λin). The difference between the energy of the transition and the intercept is the solvent contribution to the Franck−Condon barrier (λout).63,65 The λin and λout parameters are very similar for both MV derivatives 2(PF6) and 3(PF6) in accord with their almost identical structures. Simultaneous increase of the half-width of the bands conforming to eq 3 can also be noted. Therefore, it can be safely concluded that the MV complexes 2(PF6) and 3(PF6) belong to Class II and that the simplified two-level Hush model is very well obeyed. As said earlier, a “planar” delocalized ClassIII structure is computed as the most stable conformation for 3+ in the “gas phase”. Every other conformer, where πdelocalization decreases, should be thermodynamically slightly less stable. Obviously, external solvent interactions dictate the structure of the localized Class-II conformer experimentally observed. The unambiguous Class-II classification of 2(PF6) and 3(PF6) means that the electronic coupling Hab can be accurately derived from eq 7.63,64 The experimental Hab parameters (2(PF6), Hab = 262 cm−1; 3(PF6), Hab = 203 cm−1) prove quite sizable for two metal centers situated 17.7 Å (17.49 Å for computed 2+) and 19.7 Å (19.49 Å for computed 3+) apart. The significant enhancement with respect to the MV complex [Cp*(dppe)Fe−CC-B-CC−Fe(dppe)Cp*][PF6] (9(PF6), B = 2,2′-biphenyl) is very noticeable.16 Although in this compound the iron nuclei are separated by 13 bonds and a through-space distance of only 16.1 Å, the coupling parameter (Hab = 145 cm−1) is smaller than those found for 2(PF6) (−50%) and 3(PF6) (−44%). Interestingly, the Hab values for 2(PF6) and 3(PF6) for which the metal distances are close to 2 nm, almost compete with that found for compound 10(PF6) which contains the rigid bridge B = 2,7diethynylfluorene and where the metal centers are 16.1 Å apart, and the two phenyl rings are forced to be planar (Hab = 380 cm−1).26 Indeed, there are very few MV complexes with an intramolecular distance between the charge carriers approaching 2 nm.66 To the best of our knowledge, the largest electronic coupling was reported by Launay et al. for a bis-ruthenium derivative (dRuRu = 20.7 Å; Hab = 460 cm−1).67 However, it should be mentioned that the evaluation of the experimental

■

EXPERIMENTAL SECTION

General Data. Manipulations of air-sensitive compounds were performed under an argon atmosphere using standard Schlenk techniques. Tetrahydrofuran (THF) and pentane were dried and deoxygenated by distillation from sodium/benzophenone ketyl. Dichloromethane was distilled under argon from P2O5 and then from Na2CO3. Other chemicals were purchased from commercial sources and used without further purification. Infrared spectra were obtained as KBr Pellets with a Bruker IFS28 FTIR infrared spectrophotometer (4000−400 cm−1). Near-IR and UV−visible spectra were recorded as CH2Cl2 solutions, using a 1 cm long quartz cell on a Cary 5000 spectrophotometer. 1H, 13C, and 31P NMR spectra were acquired on a Bruker AVIII 400 multinuclear NMR spectrometer at ambient temperature, unless otherwise noted. Chemical shifts are reported in parts per million (δ) relative to tetramethylsilane (TMS), using the residual solvent resonances as internal references for 1H and 13 C and external H3PO4 (0.0 ppm) for 31P NMR spectra. Coupling constants (J) are reported in hertz (Hz). High-resolution mass spectra (HRMS) were recorded on a high-resolution Waters Q-Tof 2 K

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics spectrometer operating in the ESI+ mode. Cyclic voltammograms were recorded in dry CH2Cl2 solutions containing 0.1 M [Bun4N](PF6) as electrolyte, purged with argon, and maintained under argon atmosphere, using a EG&G-PAR model 362 potentiostat/galvanostat. The working electrode was a Pt disk, the counter electrode a Pt wire, and the reference electrode a saturated calomel electrode. The ferrocene/ferrocenium redox couple (E1/2 = 0.46 V) was used as an internal calibrant for the potential measurements.38 Electron paramagnetic resonance (ESR) spectra were recorded on a Bruker EMX8/2.7 (X-band) spectrometer at 66 K (liquid nitrogen under pumping). The 57Fe Mössbauer spectra were recorded with a 2.5 × 10−2 C (9.25 × 108 Bq) 57Co source using a symmetric triangular sweep mode. Computer fitting of the Mössbauer data to Lorentzian line shapes was carried out with a previously reported computer program.43 The isomer shift values are reported relative to iron foil at 298 K. Elemental analyses were conducted on a Thermo-FINIGAN Flash EA 1112 CHNS/O analyzer. Materials. Compounds Cp*(dppe)FeCl (5), Cp*(dppe)Fe−C C−C4H2S−CC−C4H2S−CCH (6), and Cp*(dppe)Fe−CC− C 4 H 2 S−CCH (7) were prepared according to published procedures.27,28 [Cp*(dppe)Fe−CC-2,5-(C4 H 2 S)-CC-2,5-(C 4H 2S)-CC−Fe(dppe)Cp*] (2). In a Schlenk flask, 0.500 g (0.60 mmol) of [Cp*(dppe)Fe−CC-2,5-(C4H2S)-CC-2′,5′-(C4H2S)-CC−H] (6), 0.248 g (0.73 mmol) of NaBPh4, and 0.416 g (0.66 mmol) of Cp*(dppe)FeCl (5) in a THF/MeOH (1/3) mixture were stirred overnight at 20 °C. Then, 0.082 g (0.73 mmol) of KOBut was added. The agitation was maintained for 1 h before the solvents were removed under reduced pressure. The solid residue was extracted with toluene (4 × 10 mL), then the solution was concentrated to 3−5, and 100 mL of pentane was added to precipitate a red brown powder. The solid was washed with pentane (3 × 10 mL) and dried in vacuo to give 0.630 g of 2 (74%). Anal. Calcd for C86H82Fe2P4S2: C, 72.98; H, 5.84; S, 4.53. Found: C, 72.68; H, 5.78; S, 4,28. HRMS (ESI+): calcd for C86H82P4S2Fe2 (M+•), 1414.35072; found, 1414.3508. FT-IR (cm−1): (Nujol), 2176 (m, CC), 2032 (s, CC); (CH2Cl2), 2150 (m, C C), 2026 (s, CC). 31P NMR (81 MHz, C6D6): δP 100.8 (s, dppe). 1 H NMR (200 MHz, C6D6, 25 °C): 7.93−6.96 (m, 42H, Ph + C4H2S/ H4); 6.50 (d, 2H, 3JHH = 3.8 Hz, C4H2S/H3); 2.51, 1.75 (2m, 8H, CH2); 1.47 (s, 30H, Cp*). 13C NMR (75 MHz, C6D6, 25 °C): 151.8 (t, 2JCP = 41 Hz, Cα); 139.6−128.0 (m, Ph); 134.4 (m, C2); 132.1 (dd, 1 JCH = 166 Hz, 2JCH = 4 Hz, C4); 124.5 (dd, 1JCH = 166 Hz, 2JCH = 4 Hz, C3); 117.0 (m, C5); 112.5 (s, Cβ); 88.3 (s, C5Me5); 86.5 (d, 3JCH = 4 Hz, CC); 31.1 (m, CH2); 10.4 (q, 1JCH = 126 Hz, C5Me5). [{Cp*(dppe)Fe−CC-2,5-C 4 H 2 S} 2 (μ-CC−CC)] (3). In a Schlenk flask, 1.295 g (1.80 mmol) of [Cp*(dppe)Fe−CC-2,5(C4H2S)-CC−H] (7), 0.326 g (1.80 mmol) of Cu(OAc)2, 40 mL of pyridine, and 274 μL of DBU (1.80 mmol) were stirred for 4 h at 50 °C. Then, the solvents were removed under reduced pressure. The solid residue was extracted with toluene (4 × 10 mL) and the resulting solution filtered on a silica pad. After concentration of the solution to 3−5 mL, 100 mL of pentane was added to precipitate a red powder. The solid was washed with pentane (3 × 10 mL) and dried in vacuo to give 1.100 g of 3 (0.764 mmol, 84%). Single crystals of 3 were grown by slow diffusion of pentane into a saturated solution of 3 in dichloromethane. Anal. Calcd for C88H82Fe2P4S2: C, 73.43; H, 5.74; S, 4.46. Found: C, 72.99; H, 5.75; S, 4.47. HRMS (ESI+): calcd for C88H82P4S2Fe2 (M+•), 1438.3507; found, 1438.3496. FT-IR (ν, cm−1): (Nujol), 2173 (w, CC), 2123 (m, CC), 2022 (s, CC); (CH2Cl2), 2124 (w, CC), 2021 (s, CC). 31P NMR (81 MHz, C6D6, 25 °C): 100.6 (s, dppe). 1H NMR (200 MHz, C6D6, 25 °C): 7.90−7.00 (m, 42H, Ph, C4H2S/H4); 6.39 (d, 1H, 3JHH = 3.48 Hz, C4H2S/H3), 2.49−1.76 (2m, 8H, CH2); 1.46 (s, 30H, Cp*). 13C NMR (50 MHz, C6D6, 25 °C): 156.9 (t, 2JCP = 38 Hz, Cα); 139.5−128.0 (m, Ph); 135.7 (m, C2); 135.2 (m, C4); 124.6 (m, C3); 115.0 (m, C5); 112.9 (s, Cβ); 88.6 (s, C5Me5); 79.1 (m, Cβ′); 77.9 (m, Cα′); 31.1 (m, CH2); 10.5 (q, 1JCH = 126 Hz, C5Me5). [{Cp*(dppe)Fe−CCCC-}2(μ-2,5-C4H2S)] (4). In a Schlenk flask, 0.065 g (0.20 mmol) of (CH3)3Si-(CC)2-2,5-(C4H2S)-(CC)2Si(CH3)3 (8), 0.300 g (0.48 mmol) of Cp*(dppe)FeCl (5), 0.066 g

(0.48 mmol) of K2CO3, 0.164 g (0.48 mmol) of NaBPh4, and 30 mL of THF are stirred at −20 °C for 15 min before to add 5 mL of MeOH. The resulting mixture is warm up to 20 °C and stirred for 16 h. Then, 0.054 g (0.48 mmol) of KOBut was added. The agitation was maintained for 1 h before the solvents were removed under reduced pressure. The solid residue was extracted with toluene (4 × 10 mL), then the solution was concentrated to 3−5 and 100 mL of pentane was added to precipitate a brown powder. The solid was washed with pentane (3 × 10 mL) and dried in vacuo to give 0.201 g of 4 (0.148 mmol, 74%). Anal. Calcd for C84H80Fe2P4S1: C, 74.34; H, 5.94; S, 2.36. Found: C, 74.23; H, 5.87; S, 4.22. HRMS (ESI+): calcd for C86H82P4S2Fe2 (M+•), 1356.3630; found, 1356.3624. FT-IR (cm−1): (Nujol), 2131 (s, CC), 1991 (s, CC); (CH2Cl2), 2121 (s, C C), 1987 (s, CC). 31P NMR (81 MHz, C6D6): 100.0 (s, dppe). 1H NMR (200 MHz, C6D6, 25 °C): 7.96−7.00 (m, 40H, Ph); 6.76 (s, 2H, C4H2S/H3,4); 2.50−1.76 (2m, 8H, CH2); 1.42 (s, 30 H, Cp*). [Cp*(dppe)Fe−CC-2,5-(C 4 H 2 S)-CC-2,5-(C4 H 2 S)-CC−Fe(dppe)Cp*] [2(PF6)]. A 0.200 g (0.141 mmol) amount of the neutral complex 2 and 0.044 g (0.134 mmol, 0.95 equiv) of ferrocenium hexafluorophosphate were combined and cooled to −80 °C before the addition of 20 mL of cold THF (−80 °C). The reaction mixture was stirred for 6 h at this temperature before adding 100 mL of cold pentane (−80 °C). Immediately, a red powder precipitated from the solution, and the mixture of solvents was removed by filtration. The solid residue was washed three times with 50 mL of cold pentane (−80 °C) affording 0.187 g of a red powder (0.187 g, 85%) upon vacuum drying. [Cp*(dppe)Fe−CC-2,5-(C4H2S)-CC−CC-2,5-(C4H2S)-C C−Fe(dppe)Cp*] [3(PF6)]. A 0.200 g (0.139 mmol) amount of the neutral complex 3 and 0.044 g (0.134 mmol, 0.95 equiv) of ferrocenium hexafluorophosphate were reacted using the recipe reported for 2(PF6). Compound 3(PF6) was isolated as a red powder (0.205 g, 93%). [Cp*(dppe)Fe−CC-2,5-(C 4 H 2 S)-CC-2,5-(C4 H 2 S)-CC−Fe(dppe)Cp*] (2(PF6)2]. A 0.200 g (0.141 mmol) amount of the neutral complex 2 and 0.091 g (0.276 mmol, 1.95 equiv) of ferrocenium hexafluorophosphate were reacted using the recipe reported for 2(PF6). Compound 2(PF6)2 was isolated as a red powder (0.206 g, 85%). [Cp*(dppe)Fe−CC-2,5-(C4H2S)-CC−CC-2,5-(C4H2S)-C C−Fe(dppe)Cp*] [3(PF6)2]. A 0.300 g (0.208 mmol) amount of the neutral complex 3 and 0.134 g (0.406 mmol, 1.95 equiv) of ferrocenium hexafluorophosphate were reacted using the recipe reported for 2(PF6). Compound 3(PF6)2 was isolated as a red powder (0.203 g, 83%). X-ray Crystal Structure Determination. A crystal of 3 was glued to a glass fiber mounted on a four-circle Nonius Kappa CCD areadetector diffractometer. Intensity data sets were collected using Mo Kα radiation (λ = 0.71073 Å) through the program COLLECT.68 Correction for the Lorentz-polarization effect, peak integration, and background determination were carried out with the program DENZO.69 Frame scaling and unit cell parameter refinement were performed with the program SCALEPACK.69 Analytical absorption corrections were performed by modeling the crystal faces using NUMABS.70 The structure was solved with the triclinic space group P1. Iron and phosphorus atoms were located using the direct methods with the program SIR97.71 The complete model was found from successive Fourier calculations using SHELXL-97.72 A summary of the details about crystal data, collection parameters, and refinement is documented in Table S1 (Supporting Information), and additional crystallographic details are provided in the CIF file. Computational Details. DFT calculations were carried out using the Amsterdam Density Functional (ADF) program.73,74 Electron correlation was treated within the local density approximation (LDA) in the Vosko−Wilk−Nusair parametrization.75 Nonlocal corrections were added to the exchange and correlation energies using the BP8676 and B3LYP77 functionals. Calculations were performed using the standard ADF triple-ζ quality basis set. Full geometry optimizations (assuming C1 symmetry) were carried out on each complex, using the analytical gradient method implemented by Versluis and Ziegler.78 L

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics The nature of the stationary points after optimization was checked by calculations of the harmonic vibrational frequencies. Computed EPR properties were accomplished using the ESR procedure developed by van Lenthe and co-workers.79 The g-tensor components were obtained using self-consistent spin-unrestricted DFT calculations after incorporating the relativistic spin−orbit coupling by first order perturbation theory from a ZORA Hamiltonian, using the BP86 functional for nonlocal corrections to the exchange and correlation energies.80 TD-DFT calculations80 were performed on the optimized structures using the long-range CAMY-B3LYP functional81 and taking into account the solvation effects using the conductor-like screening model (COSMO)82 with a dielectric constant simulating dichloromethane solvent. Molecular orbitals and spin densities were plotted with the ADF-GUI package.74 Full details of the structure determination have also been deposited with the Cambridge Crystallographic Data Centre as CCDC 926490. Copies of this information may be obtained free of charge from The Director, CCDC, 12 Union Street, Cambridge CB2 1EZ, U.K. (fax, + 44-1223-336-033; e-mail, [email protected].

■

(8) (a) Zhao, X.; Huang, C.-C.; Gulcur, M.; Batsanov, A. S.; Baghernejad, M.; Hong, W.; Bryce, M. R.; Wandlowski, T. Chem. Mater. 2013, 25, 4340−4347. (b) Prins, F.; Monrabal-Capilla, M.; Osorio, E. A.; Coronado, E.; Van der Zant, H. S. J. Adv. Mater. 2011, 23, 1545−1549. (c) Low, P. J. Dalton Trans. 2005, 2821−2824. (d) Wuttke, E.; Hervault, Y.-M.; Polit, W.; Linseis, M.; Erler, P.; Rigaut, S.; Winter, R. F. Organometallics 2014, 33, 4672−4686. (e) O’Hanlon, D. C.; Cohen, B. W.; Moravec, D. B.; Dallinger, R. F.; Hopkins, M. D. J. Am. Chem. Soc. 2014, 136, 3127−3136. (f) Zhu, H.; Pookpanratana, S. J.; Bonevich, J. E.; Natoli, S. N.; Hacker, C. A.; Ren, T.; Suehle, J. S.; Richter, C. A.; Li, Q. ACS Appl. Mater. Interfaces 2015, 7, 27306−27313. (9) Lissel, F.; Schwarz, F.; Blacque, O.; Riel, H.; Lörtscher, E.; Venkatesan, K.; Berke, H. J. Am. Chem. Soc. 2014, 136, 14560−14569. (10) Burgun, A.; Gendron, F.; Sumby, C.; Roisnel, T.; Cador, O.; Costuas, K.; Halet, J.-F.; Bruce, M. I.; Lapinte, C. Organometallics 2014, 33, 2613−2627. (11) (a) Lu, Y.; Quardokus, R.; Lent, C. S.; Justaud, F.; Lapinte, C.; Kandel, S. A. J. Am. Chem. Soc. 2010, 132, 13519−13524. (b) Quardokus, R. C.; Lu, Y.; Wasio, N. A.; Lent, C. S.; Justaud, F.; Lapinte, C.; Kandel, S. A. J. Am. Chem. Soc. 2012, 134, 1710−1714. (c) Wasio, N. A.; Quardokus, R. C.; Forrest, R. P.; Corcelli, S. A.; Lu, Y.; Lent, C. S.; Justaud, F.; Lapinte, C.; Kandel, S. A. J. Phys. Chem. C 2012, 116, 25486−25492. (12) Burgun, A.; Gendron, F.; Schauer, P. A.; Skelton, B. W.; Low, P. J.; Costuas, K.; Halet, J.-F.; Bruce, M. I.; Lapinte, C. Organometallics 2013, 32, 5015−5025. (13) Frampton, M. J.; Anderson, H. L. Angew. Chem., Int. Ed. 2007, 46, 1028−1064. (14) Weisbach, N.; Baranova, Z.; Gauthier, S.; Reibenspies, J. H.; Gladysz, J. A. Chem. Commun. 2012, 48, 7562−7564. (15) Le Narvor, N.; Lapinte, C. Organometallics 1995, 14, 634−639. (16) Ghazala, S. I.; Paul, F.; Toupet, L.; Roisnel, T.; Hapiot, P.; Lapinte, C. J. Am. Chem. Soc. 2006, 128, 2463−2476. (17) (a) Tanaka, Y.; Shaw-Taberlet, J. A.; Justaud, F.; Cador, O.; Roisnel, T.; Akita, M.; Hamon, J.-R.; Lapinte, C. Organometallics 2009, 28, 4656−4669. (b) Shaw-Taberlet, J. A.; Sinbandit, S.; Roisnel, T.; Hamon, J.-R.; Lapinte, C. Organometallics 2006, 25, 5311−5325. (18) de Montigny, F.; Argouarch, G.; Costuas, K.; Halet, J.-F.; Roisnel, T.; Toupet, L.; Lapinte, C. Organometallics 2005, 24, 4558− 4572. (19) Le Stang, S.; Paul, F.; Lapinte, C. Organometallics 2000, 19, 1035−1043. (20) Stahl, J.; Bohling, J. C.; Bauer, E. B.; Peters, T. B.; Mohr, W.; Martin-Alvarez, J. M.; Hamplel, F.; Gladysz, J. A. Angew. Chem., Int. Ed. 2002, 41, 1871−1876. (21) Fox, M. A.; Le Guennic, B.; Roberts, R. L.; Brue, D. A.; Yufit, D. S.; Howard, J. A. K.; Manca, G.; Halet, J.-F.; Hartl, F.; Low, P. J. J. Am. Chem. Soc. 2011, 133, 18433−18446. (22) (a) Bruce, M. I.; Le Guennic, B.; Scoleri, N.; Zaitseva, N. N.; Halet, J.-F. Organometallics 2012, 31, 4701−4706. (b) Lohan, M.; Justaud, F.; Roisnel, T.; Ecorchard, P.; Lang, H.; Lapinte, C. Organometallics 2010, 29, 4804−4817. (23) Gendron, F.; Burgun, A.; Skelton, B. W.; White, A. H.; Roisnel, T.; Bruce, M. I.; Halet, J.-F.; Lapinte, C.; Costuas, K. Organometallics 2012, 31, 6796−6811. (24) Bruce, M. I.; Costuas, K.; Gendron, F.; Halet, J.-F.; Jevric, M.; Skelton, B. W. Organometallics 2012, 31, 6555−6566. (25) Burgun, A.; Gendron, F.; Roisnel, T.; Sinbandhit, S.; Costuas, K.; Halet, J.-F.; Bruce, M. I.; Lapinte, C. Organometallics 2013, 32, 1866−1875. (26) Malvolti, F.; Rouxel, C.; Mongin, O.; Hapiot, P.; Toupet, L.; Blanchard-Desce, M.; Paul, P. Dalton Trans. 2011, 40, 6616−6618. (27) Roger, C.; Hamon, P.; Toupet, L.; Rabaâ, H.; Saillard, J.-Y.; Hamon, J.-R.; Lapinte, C. Organometallics 1991, 10, 1045−1054. (28) Roué, S.; Lapinte, C. J. Organomet. Chem. 2005, 690, 594−604. (29) (a) Brandsma, L.; Verkruijsse, H. D. Preparative Polar Organic Chemistry; Springer Verlag: Berlin, Germany, 1987; Vol. 1.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.organomet.6b00209. Crystallographic data, details of data collection, and structure refinement parameters for 3 (Table S1), computed and experimental ESR parameters for 2+ and 3+ (Table S2), 57Fe Mössbauer spectrum for 2(PF6) at 80 K (Figure S1), UV−vis spectra for 2−4 and 6 (Figure S2), UV−vis spectra for 3(PF6)n (n = 0−2) (Figure S3), and Mulliken atomic spin densities for 2+ (Figure S4) (PDF) Cartesian coordinates for all calculated optimized geometries (MOL) CIF file giving crystallographic data for 3 (CIF)

■

AUTHOR INFORMATION

Corresponding Authors

*(J.-H.F.) E-mail: [email protected]. *(C.L.) E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■ ■

ACKNOWLEDGMENTS This article is dedicated to our good friend and colleague JeanRené Hamon on the occasion of his 60th birthday. REFERENCES

(1) Paul, F.; Lapinte, C. Coord. Chem. Rev. 1998, 178−180, 431−509. (2) (a) Aguirre-Etcheverry, P.; O’Hare, D. Chem. Rev. 2010, 110, 4839−4864. (b) Costuas, K.; Rigaut, S. Dalton Trans. 2011, 40, 5643− 5658. Low, P. J. Coord. Chem. Rev. 2013, 257, 1507−1532. (3) Halet, J.-F.; Lapinte, C. Coord. Chem. Rev. 2013, 257, 1584−1613. (4) Le Narvor, N.; Toupet, L.; Lapinte, C. J. Am. Chem. Soc. 1995, 117, 7129−7138. (5) (a) Dembinski, R.; Bartik, T.; Bartik, B.; Jaeger, M.; Gladysz, J. A. J. Am. Chem. Soc. 2000, 122, 810−822. (b) Bruce, M. I.; Low, P. J.; Costuas, K.; Halet, J.-F.; Best, S. P.; Heath, G. A. J. Am. Chem. Soc. 2000, 122, 1949−1962. (c) Guillemot, M.; Toupet, L.; Lapinte, C. Organometallics 1998, 17, 1928−1930. (6) Bruce, M. I.; Costuas, K.; Davin, T.; Ellis, B. E.; Halet, J.-F.; Lapinte, C.; Low, P. J.; Smith, K. M.; Skelton, B. W.; Toupet, L.; White, A. H. Organometallics 2005, 24, 3864−3881. (7) Coat, F.; Lapinte, C. Organometallics 1996, 15, 477−480. M

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX

Article

Organometallics (b) Brandsma, L. Preparative Acetylenic Chemistry; Elsevier: Amsterdam, The Netherlands, 1988. (30) Weyland, T.; Lapinte, C.; Frapper, G.; Calhorda, M. J.; Halet, J.F.; Toupet, L. Organometallics 1997, 16, 2024−2031. (31) Coat, F.; Paul, F.; Lapinte, C.; Toupet, L.; Costuas, K.; Halet, J.F. J. Organomet. Chem. 2003, 683, 368−378. (32) Schermann, G.; Grösser, T.; Hample, F.; Hirsch, A. Chem.Eur. J. 1997, 3, 1105−1112. (33) Hamon, P.; Toupet, L.; Hamon, J.-R.; Lapinte, C. Organometallics 1996, 15, 10−12. (34) (a) Zhu, Y.; Millet, D. B.; Wolf, M. O.; Rettig, S. J. Organometallics 1999, 18, 1930−1938. (b) Lewis, J.; Long, N. J.; Raithby, P. R.; Shields, G. P.; Wong, W.-Y.; Younus, M. J. Chem. Soc., Dalton Trans. 1997, 4283−4288. (c) Antolini, L.; Tedesco, E.; barbarella, G.; Favaretto, L.; Sotgiu, L.; Zambianchi, M.; Casarini, D.; Gigli, G.; Cingolani, R. J. Am. Chem. Soc. 2000, 122, 9006−9013. (35) Hamon, P.; Justaud, F.; Cador, O.; Hapiot, P.; Rigaut, S.; Toupet, L.; Ouahab, L.; Stueger, H.; Hamon, J.-R.; Lapinte, C. J. Am. Chem. Soc. 2008, 130, 17372−17383. (36) (a) Myers, R. L.; Shain, I. Anal. Chem. 1969, 41, 980−980. Hapiot, P.; Kispert, L. D.; Konovalov, V. V.; Savéant, J.-M. J. Am. Chem. Soc. 2001, 123, 6669−6667. (b) Andrieux, C. P.; Savéant, J.-M. J. Electroanal. Chem. Interfacial Electrochem. 1974, 57, 27−33. (c) Guerro, M.; Carlier, R.; Boubekeur, K.; Lorcy, D.; Hapiot, P. J. Am. Chem. Soc. 2003, 125, 3159−3167. (37) Ribou, A.-C.; Launay, J.-P.; Takahashi, K.; Nihara, T.; Tarutani, S.; Spangler, C. W. Inorg. Chem. 1994, 33, 1325−1329. (38) Connelly, N. G.; Geiger, W. E. Chem. Rev. 1996, 96, 877−910. (39) Levitus, M.; Schmieder, K.; Ricks, H.; Shimizu, K. D.; Bunz, U. H. F.; Garcia-Garibay, M. A. J. Am. Chem. Soc. 2001, 123, 4259−4265. (40) Paul, F.; Mevellec, J.-Y.; Lapinte, C. Dalton Trans. 2002, 1783− 1790. (41) Denis, R.; Toupet, L.; Paul, F.; Lapinte, C. Organometallics 2000, 19, 4240−4251. (42) Paul, F.; Toupet, L.; Thépot, J.-Y.; Costuas, K.; Halet, J.-F.; Lapinte, C. Organometallics 2005, 24, 5464−5478. (43) Varret, F.; Mariot, J.-P.; Hamon, J.-R.; Astruc, D. Hyperfine Interact. 1988, 39, 67−81. (44) (a) Argouarch, G.; Thominot, P.; Paul, F.; Toupet, L.; Lapinte, C. C. R. Chim. 2003, 6, 209−222. (b) Guillaume, V.; Thominot, P.; Coat, F.; Mari, A.; Lapinte, C. J. Organomet. Chem. 1998, 565, 75−80. (45) Guillaume, V.; Mahias, V.; Mari, A.; Lapinte, C. Organometallics 2000, 19, 1422−1426. (46) Costuas, K.; Cador, O.; Justaud, F.; Le Stang, S.; Paul, F.; Monari, A.; Evangelisti, S.; Toupet, L.; Lapinte, C.; Halet, J.-F. Inorg. Chem. 2011, 50, 12601−12622. (47) Makhoul, R.; Sahnoune, H.; Dorcet, V.; Halet, J.-F.; Hamon, J.R.; Lapinte, C. Organometallics 2015, 34, 3314−3326. (48) (a) Dong, T.-Y.; Hendrickson, D. N.; Pierpont, C. G.; Moore, M. F. J. Am. Chem. Soc. 1986, 108, 963−971. (b) Dong, T.-Y.; Sohel, C.-C.; Hwang, M.-Y.; Lee, T. Y.; Yeh, S.-K.; Wen, Y.-S. Organometallics 1992, 11, 573−582. (49) Roué, S.; Le Stang, S.; Toupet, L.; Lapinte, C. C. R. Chim. 2003, 6, 353−366. (50) Bertini, I.; Luchinat, C. NMR of Paramagnetic Molecules in Biological Systems; The Benjamin/Cummings Publishing Company, Inc.: Menlo Park, CA, 1986. (51) Wicholas, M.; Mustacich, R.; Jayne, D. J. Am. Chem. Soc. 1972, 94, 4518−4522. (52) La Mar, G. N.; Eaton, G. R.; Holm, R. H.; Walker, F. A. J. Am. Chem. Soc. 1973, 95, 63−75. (53) Weyland, T.; Costuas, K.; Mari, A.; Halet, J.-F.; Lapinte, C. Organometallics 1998, 17, 5569−5579. (54) Kahn, O. Molecular Magnetism; VCH Publishers: New York, 1993. (55) Paul, F.; Bondon, A.; Da Costa, G.; Malvoti, F.; Sinbandit, S.; Cador, O.; Costuas, K.; Toupet, L.; Boillot, M. L. Inorg. Chem. 2009, 48, 10608−10624. (56) Lapinte, C. J. Organomet. Chem. 2008, 693, 793−801.

(57) Ruiz, E.; Cano, J.; Alvarez, S.; Alemany, P. J. Comput. Chem. 1999, 20, 1391−1400. (58) Costuas, K.; Paul, F.; Toupet, L.; Halet, J.-F.; Lapinte, C. Organometallics 2004, 23, 2053−2068. (59) Robin, M. B.; Day, P. Adv. Inorg. Chem. Radiochem. 1968, 10, 247−422. (60) Hush, N. S. Trans. Faraday Soc. 1961, 57, 557−580. Hush, N. S. Prog. Inorg. Chem. 1967, 8, 391−444. (61) Demadis, K. D.; Hartshorn, C. M.; Meyer, T. J. Chem. Rev. 2001, 101, 2655−2686. (62) Crutchley, R. J. Adv. Inorg. Chem. 1994, 41, 273−325. (63) Astruc, D. Electron Transfer and Radical Processes in TransitionMetal Chemistry; VCH: New York, 1995. (64) Sauvage, J.-P.; Collin, J.-P.; Chambron, J.-C.; Guillerez, S.; Coudret, C. Chem. Rev. 1994, 94, 993−1019. (65) Tom, G. M.; Creutz, C.; Taube, H. J. Am. Chem. Soc. 1974, 96, 7827−7829. (66) Cao, Z.; Xi, B.; Jodoin, D. S.; Zhang, L.-Y.; Cummings, S. P.; Gao, Y.; Tyler, S. F.; Fanwick, P. E.; Crutchley, R. J.; Ren, T. J. Am. Chem. Soc. 2014, 136, 12174−12183. (67) Fraysse, S.; Coudret, C.; Launay, J.-P. J. Am. Chem. Soc. 2003, 125, 5880−5886. (68) Nonius Kappa CCD Software; Nonius BV: Delft, The Netherlands, 1999. (69) Otwinowski, Z.; Minor, W. Processing of X-ray Diffraction Data Collected in Oscillation Mode. In Methods in Enzymology, Macromolecular Crystallography, Part A; Carter, C. W., Sweet, R. M., Eds.; Academic Press: New York, 1997; Vol. 276, pp 307−326. (70) Coppens, P. In Crystallographic Computing; Munksgaard Publishers: Copenhagen, Denmark, 1970. (71) Altomare, A.; Burla, M. C.; Camali, M.; Cascarano, G.; Giacovazzo, C.; Guagliardi, A.; Moliterni, A. G. G.; Polidori, G.; Spagna, R. J. Appl. Crystallogr. 1999, 32, 115−119. (72) Sheldrick, G. M. Acta Crystallogr., Sect. A: Found. Crystallogr. 2008, 64, 112−122. (73) (a) Te Velde, G.; Bickelhaupt, F. M.; van Gisbergen, S. J. A.; Fonseca Guerra, C.; Baerends, E. J.; Snijders, J. G.; Ziegler, T. J. J. Comput. Chem. 2001, 22, 931−967. (b) Fonseca Guerra, C.; Snijders, J. G.; te Velde, G.; Baerends, E. J. Theor. Chem. Acc. 1998, 99, 391− 403. (74) ADF2014.07; SCM, Theoretical Chemistry; Vrije Universiteit: Amsterdam, The Netherlands. http://www.scm.com. (75) Vosko, S. H.; Wilk, L.; Nusair, M. Can. J. Phys. 1980, 58, 1200− 1211. (76) (a) Becke, A. D. Phys. Rev. A: At., Mol., Opt. Phys. 1988, 38, 3098−3100. (b) Perdew, J. P. Phys. Rev. B: Condens. Matter Mater. Phys. 1986, 33, 8822−8824. (77) Becke, A. D. J. Chem. Phys. 1993, 98, 5648−5652. (78) Verluis, L.; Ziegler, T. J. Chem. Phys. 1988, 88, 322−328. (79) (a) van Lenthe, E.; van der Avoird, A.; Wormer, P. E. S. J. Chem. Phys. 1998, 108, 4783−4796. (b) van Lenthe, E.; van der Avoird, A.; Wormer, P. E. S. J. Chem. Phys. 1997, 107, 2488−2408. (80) (a) van Gisbergen, S. J. A.; Snijders, J. G.; Baerends, E. J. Comput. Phys. Commun. 1999, 118, 119−138. (b) van Lenthe, E.; Baerends, E. J. J. Comput. Chem. 2003, 24, 1142−1156. (81) Seth, M.; Ziegler, T. J. Chem. Theory Comput. 2012, 8, 901−907. (82) Klamt, A. J. Phys. Chem. 1995, 99, 2224−2235.

N

DOI: 10.1021/acs.organomet.6b00209 Organometallics XXXX, XXX, XXX−XXX