Electronic and Structural Effects of Inner Sphere Coordination of

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Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

Electronic and Structural Effects of Inner Sphere Coordination of Chloride to a Homoleptic Copper(II) Diimine Complex Valentina Leandri,† Quentin Daniel,‡ Hong Chen,‡ Licheng Sun,‡,§ James M. Gardner,† and Lars Kloo*,† †

Applied Physical Chemistry, Department of Chemistry, KTH Royal Institute of Technology, SE-10044, Stockholm, Sweden Organic Chemistry, Centre of Molecular Devices, Department of Chemistry, School of Chemical Science and Engineering, KTH Royal Institute of Technology, SE-100 44, Stockholm, Sweden § State Key Laboratory of Fine Chemicals, DUT-KTH Joint Research Center on Molecular Devices, Dalian University of Technology (DUT), 116024 Dalian, China ‡

S Supporting Information *

ABSTRACT: The reaction of CuCl2 with 2,9-dimethyl-1,10phenanthroline (dmp) does not lead to the formation of [Cu(dmp)2](Cl)2 but instead to [Cu(dmp)2Cl]Cl, a 5-coordinated complex, in which one chloride is directly coordinated to the metal center. Attempts at removing the coordinated chloride by changing the counterion by metathesis were unsuccessful and resulted only in the exchange of the noncoordinated chloride, as confirmed from a crystal structure analysis. Complex [Cu(dmp)2Cl]PF6 exhibits a reversible cyclic voltammogram characterized by a significant peak splitting between the reductive and oxidative waves (0.85 and 0.60 V vs NHE, respectively), with a half-wave potential E1/2 = 0.73 V vs NHE. When reduced electrochemically, the complex does not convert into [Cu(dmp)2]+, as one may expect. Instead, [Cu(dmp)2]+ is isolated as a product when the reduction of [Cu(dmp)2Cl]PF6 is performed with L-ascorbic acid, as confirmed by electrochemistry, NMR spectroscopy, and diffractometry. [Cu(dmp)2]2+ complexes can be synthesized starting from Cu(II) salts with weakly and noncoordinating counterions, such as perchlorate. Growth of [Cu(dmp)2](ClO4)2 crystals in acetonitrile results in a 5coordinated complex, [Cu(dmp)2(CH3CN)](ClO4)2, in which a solvent molecule is coordinated to the metal center. However, solvent coordination is associated with a dynamic decoordination−coordination behavior upon reduction and oxidation. Hence, the cyclic voltammogram of [Cu(dmp)2(CH3CN)]2+ is identical to the one of [Cu(dmp)2]+, if the measurements are performed in acetonitrile. The current results show that halide ions in precursors to Cu(II) metal−organic coordination compound synthesis, and most likely also other multivalent coordination centers, are not readily exchanged when exposed to presumed strongly binding and chelating ligand, and thus special care needs to be taken with respect to product characterization.

1. INTRODUCTION

which causes the oxidation of copper(I) to copper(II) to be thermodynamically less favorable.5,6 For this reason, the redox potential for the couple [Cu(dmp)2]2+/+ (dmp = 2,9-dimethyl1,10-phenanthroline) is substantially more positive than that of [Cu(phen)2]2+/+ (phen = 1,10-phenanthroline).7 In 2005, Hattori et al. reported the use of [Cu(dmp)2]2+/+ and [Cu(phen)2]2+/+ as redox mediators in dye sensitized solar cells (DSSCs).8 Their study showed significantly higher performances for the [Cu(dmp)2]2+/+ couple over [Cu(phen)2]2+/+. The differences in performance are largely attributed to the more positive redox potential of [Cu(dmp)2]2+/+, which resulted in devices displaying higher open-circuit voltage (Voc) and power conversion efficiency. The maximum conversion efficiency value recorded was 2.2%

Transition metal complexes with chelating polyimines of the 1,10-phenanthroline family have been studied in detail since the early 1940s.1 In particular, copper(I) complexes prepared from 2,9-disubstituted-1,10-phenanthroline ligands have received particular attention due to their interesting photophysical, chemical, and structural properties.2,3 Complexes of Cu(I) adopt a tetrahedral or pseudo-tetrahedral geometry and are often bright orange or red due to MLCT (d−π*) electronic transitions. In the absence of restricting steric effects, these complexes may be readily oxidized to the more stable squareplanar, often green, Cu(II) species.4 The geometrical changes associated with the copper oxidation states have a remarkable impact on the electrochemical properties of complexes derived from 1,10-phenanthroline with substituents in the 2,9-positions. In the latter case, bulky substituents cause the geometry of Cu(II) to be nonplanar by inducing a strain in the molecule, © XXXX American Chemical Society

Received: February 5, 2018

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DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

NMR (acetone d6): δ = 8.75 (d, J = 8.21 Hz, 1H), 8.23 (s, 1H), 7.98 (d, J = 8.24 Hz, 1H), 2.52 (s, 3H) ppm. Single crystals of [Cu(dmp)2]PF6 suitable for X-ray analysis were obtained by slow evaporation of a dichloromethane solution containing [Cu(dmp)2]PF6. The crystals obtained were red and needle shaped. Electrochemical Characterization. Electrochemical experiments were performed in a three-electrode electrochemical cell with a platinum wire as counter electrode, Ag/AgNO3 (10 mM/acetonitrile) as reference electrode, and a glassy carbon disk (Ø = 3 mm) as the working electrode. The supporting electrolyte used was a 0.1 M (nBu)4NPF6/acetonitrile solution. The measurements were performed using an Ivium Technologies vertex potentiostat. Ferrocene (Fc) was used as internal standard and the redox potentials reported considering Fc+/Fc in acetonitrile solution to be 0.63 V vs NHE. Single Crystal X-ray Diffraction. Crystals suitable for single crystal X-ray diffraction experiments were grown in different experimental conditions, as mentioned in the Synthesis of Metal Complexes section. The data were collected at room temperature around 298 K on a Bruker Adventure D8 diffractometer employing Mo K alpha radiation (λ = 0.71073 Å). The data reduction was carried with the Apex 3 software, and multiscan absorption14 was used on the original diffraction data. The crystal structure was solved by using direct methods with the software SHELXS, and the obtained structural models were further refined with the software SHELXL.15 Anisotropic thermal parameters were refined for the non-hydrogen atoms. Hydrogen atoms were added using a riding model. Final atomic positional coordinates, with estimated standard deviations, bond lengths, and angles, have been deposited at the Cambridge Crystallographic Data Centre and were allocated deposition numbers CCDC 1819088, 1819090, and 1822055.

for DSSCs based on [Cu(dmp)2]2+/+ under weak solar light irradiation, opening wide possibilities for further enhancement. More recently, several studies have reported highly efficient DSSCs (>7% power conversion efficiency) based on [Cu(dmp)2]2+/+ in combination with organic dyes.9−13 Unlike Hattori et al., many of these studies report the synthesis of [Cu(dmp)2]2+/+ starting from CuCl2. Here, we show that the decision of starting the synthesis from copper(II) chloride leads to a different compound with different electrochemical properties, potentially generating undesired effects in the final electrolyte solution.

2. EXPERIMENTAL SECTION General Information. All chemicals were purchased from SigmaAldrich and used as received unless noted otherwise. NMR spectra were recorded with an instrument operating at 400 MHz. Synthesis of Metal Complexes. Synthesis of complex 2 [Cu(dmp)2Cl]Cl: CuCl2 (250 mg, 1.86 mmol) was dissolved in ethanol (120 mL). After complete dissolution, 2.2 equiv of 2,9dimethyl-1,10-phenanthroline (852 mg, 4.09 mmol) was added as solid, and the solution and after a few minutes, a green precipitate started to form. The suspension was left to fully react by stirring at room temperature for 2 h. After this time, the green precipitate was collected by vacuum filtration using a fritted funnel. The solid was washed with cold ethanol, diethyl ether, and dried under vacuum (75% yield; 769 mg, 1.39 mmol). Synthesis of complex 4 [Cu(dmp)2Cl]PF6: [Cu(dmp)2Cl]Cl (150 mg, 0.272 mmol) was dissolved in a 1:2 ethanol/water mixture (volume ratio, 50 mL). To this solution, 20 equiv of ammonium hexafluorophosphate was added (887 mg, 5.44 mmol). After a few minutes, a green precipitate started to form. The suspension was left stirring for 1 h at room temperature. The green precipitate was then collected by vacuum filtration with a fritted funnel, washed thoroughly with water, diethyl ether, and dried under reduced pressure (90% yield; 162 mg, 0.245 mmol). Single crystals of [Cu(dmp)2Cl]PF6 suitable for X-ray analysis were obtained by slow evaporation of an acetonitrile solution containing [Cu(dmp)2Cl]PF6. The crystals obtained were green and squared shaped. UV−vis of complex 4, and of the complex obtained from the reduction of complex 4 via bulk electrolysis, are shown in the Supporting Information (Figure S3). Synthesis of complex 5 [Cu(dmp)2](ClO4)2: Cu(ClO4)2·6H2O (300 mg, 0.810 mmol) was dissolved in ethanol (50 mL). After complete dissolution, 2.2 equiv of 2,9-dimethyl-1,10-phenanthroline (371 mg, 1.78 mmol) was added as solid, affording a green precipitate after a few minutes. The suspension was allowed to fully react by stirring at room temperature for 2 h. [Cu(dmp)2](ClO4)2 (5) was collected as a green solid by vacuum filtration, washed with cold ethanol, diethyl ether, and dried under vacuum. When dried, the color of the solid turns from green to purple (80% yield; 311 mg, 0.648 mmol). Single crystals of [Cu(dmp)2(CH3CN)](ClO4)2·2(CH3CN) (6) suitable for X-ray analysis were obtained by slow evaporation of an acetonitrile solution containing [Cu(dmp)2](ClO4)2. The crystals obtained were green and squared shaped. Caution! Although no accident occurred during the experimental work associated with this article, it should be pointed out that perchlorates are hazardous, as strong oxidants, and can in combination with substances that can be oxidized cause explosive reactions. Synthesis of complex 7 [Cu(dmp)2]PF6: Complex 4 (200 mg, 0.302 mmol) was dissolved in 50 mL of acetonitrile. To this solution, 10 equi of L-ascorbic acid was added (532 mg, 3.02 mmol). Upon addition of L-ascorbic acid, the solution quickly turns from green to orange-red. The mixture was further stirred at room temperature for 1 h. The excess of L-ascorbic acid was removed by filtration, and the solution was dried under rotary evaporation. The resulting red powder was dispersed in 100 mL of deionized water and sonicated to remove any traces of L-ascorbic acid. The resulting bright orange-red solid was then collected by vacuum filtration, washed thoroughly with water, diethyl ether, and dried under vacuum (95% yield; 179 mg, 0.287 mmol). 1H

3. RESULTS AND DISCUSSION Due to the well-known explosive nature of perchlorate salts,16 a cheap and readily available alternative to many inorganic salts are their halides. Copper halides salts are no exception, and they have been widely used as starting materials for the synthesis of a variety of copper complexes.17−20 Scheme 1 and Table 1 report the molecular structures and the crystallographic details, respectively, of the copper complexes discussed in this work. Scheme 1. Molecular Structures of Complexes [Cu(dmp)2Cl]Cl (2), [Cu(dmp)2Cl]PF6 (4), [Cu(dmp)2(CH3CN)](ClO4)2 (6), and [Cu(dmp)2]PF6 (7)

More recently, copper chloride salts have been used for the synthesis of homoleptic compounds based on substituted bipyridine and phenanthroline ligands.10,11 Scheme 2 shows both the synthetic path claimed to yield the expected product [Cu(dmp)2]2+ (1), and the correct complex obtained from such synthesis (2). When employing a copper(II) salt with a strongly coordinating counterion, such as Cl−, a different structure B

DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

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Table 1. Selected Bond Distances (Å) and Angles (deg) for [Cu(dmp)2Cl]PF6 (4), [Cu(dmp)2(CH3CN)](ClO4)2 (6), and [Cu(dmp)2]PF6 (7) [Cu(dmp)2Cl]PF6 (4) Cu−N1 Cu−N2 Cu−N3 Cu−N4 Cu−Cl Cl−Cu−N1 Cl−Cu−N2 Cl−Cu−N3 Cl−Cu−N4 N1−Cu−N2 N1−Cu−N3 N1−Cu−N4 N2−Cu−N3 N2−Cu−N4 N3−Cu−N4

1.994(8) 2.090(8) 2.006(8) 2.151(9) 2.316(3) 89.1(3) 133.6(2) 85.9(2) 118.2(3) 80.0(3) 174.3(4) 104.5(3) 101.5(3) 108.2(3) 80.3(4)

[Cu(dmp)2(CH3CN)](ClO4)2 (6) Cu−N1 Cu−N2 Cu−N3 Cu−N4 Cu−N5 N5−Cu−N1 N5−Cu−N2 N5−Cu−N3 N5−Cu−N4 N1−Cu−N2 N1−Cu−N3 N1−Cu−N4 N2−Cu−N3 N2−Cu−N4 N3−Cu−N4

2.031(3) 2.057(5) 2.021(3) 2.161(4) 2.040(6) 88.2(2) 147.0(2) 83.7(2) 105.6(2) 81.4(2) 166.6(2) 112.4(1) 99.8(2) 107.3(1) 80.1(1)

Scheme 2. Synthesis of [Cu(dmp)2Cl]Cla

[Cu(dmp)2]PF6 (7) Cu−N1 Cu−N2 Cu−N3 Cu−N4

2.032(4) 2.068(5) 2.061(6) 2.052(6)

N1−Cu−N2 N1−Cu−N3 N1−Cu−N4 N2−Cu−N3 N2−Cu−N4 N3−Cu−N4

82.0(2) 131.0(2) 127.7(2) 113.2(2) 125.8(2) 82.3(2)

Scheme 3. Synthesis of [Cu(dmp)2Cl]PF6a

a

Reagent and conditions: (ii) NH4PF6 (excess), 1:2 = ethanol:water (volume ratio), RT, 1 h.

a

Reagent and conditions: (i) 2,9-dimethyl-1,10-phenanthroline (2.2 equiv), ethanol, RT, 2 h.

than the expected four-coordinated complex 1 is obtained.21 If we indicate any strongly coordinating counterion with the letter X and a ligand with the letter L, the resulting Cu(II) complex can be described with the general formula [Cu(L)2X]X. The unit [Cu(L)2X] is a coordination complex with an overall +1 charge and distances between the copper center and coordinated X− ions typical for direct coordination. This class of Cu(II) complexes is known since the early 80’s, and many compounds in which the ligand is a derivative of bipyridine or 1,10-phenanthroline have been reported.22−24 Furthermore, attempts to exchange the counterion of compound 2 do not affect the chloride that is directly coordinated to the metal center, but instead replacement of the noncoordinated Cl− ion is observed (Scheme 3). Complex 4 was isolated as a green powder, and its structure was determined by single crystal X-ray diffraction (Figure 1). The difficulty of removing the coordinated chloride is related to the relatively strong interaction between the metal center and the chloride atom. Indeed, Cu(II) complexes have a d9 configuration which promotes a pseudo-Jahn−Teller (PJT) structural distortion in symmetries rendering degeneration of frontier molecular orbitals dominated by the metal center dorbitals.25−27 As a result, they tend to adopt a more flattened configuration than the corresponding Cu(I) analogues, leaving the metal center more exposed to further coordination. For this reason, Cu(II) complexes can be often found in a five-

Figure 1. (a) Molecular structure of the cation of complex 4: [Cu(dmp)2Cl]+. Ellipsoids shown at 35% probability; hydrogen atoms are omitted for the sake of clarity. (b) Inner coordination sphere of the Cu(II) metal center of complex 4. The metal center lies in a plane defined by the ligand atoms N2, N4, and Cl.

C

DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry coordinate configuration and, despite the lack of evidence, even the existence of six-coordinated structures have been postulated.23,28 The structure of [Cu(dmp)2Cl]PF6 (Figure 1) is similar to those previously reported for [Cu(dmp) 2 (H 2 O)](CF 3 SO 3 ) 2 and [Cu(dmbp) 2 (H 2 O)](ClO4)2.28,29 The metal center can be described as a distorted trigonal bipyramidal geometry, in which the coordinated chloride is an equatorial ligand (Figure 1b), with a pseudoD3h coordination center. Selected bond distances and angles for [Cu(dmp)2Cl]PF6 are listed in Table 1. The Cu−Cl distance in complex 4 (2.316(3) Å) is typical of a bond length between a Cu(II) center and a Cl− ligand in an axial position.23,22,30 Compared to [Cu(dmp)2(H2O)](CF3SO3)2, the Cu−N bond distances in the two compounds are quite similar, while the Cu−Cl distance is significantly longer than the Cu−O distance (2.075 Å) due to the larger radius of Cl− compared to oxygen in the water ligand.28 The dihedral angle between the two planes defined by the phenanthroline ligands is 59.1°, which is similar to that of the planes defined by the phenanthroline ligands in a four-coordinated [Cu(dmp)2]2+ complex (59.4°).29 Solvent coordination at the metal center can be avoided by employing noncoordinating solvents.29 Performing the synthesis from a copper(II) salt of a noncoordinating counterion (Y), such as Cu(ClO4)2 or Cu(BF4)2,31 leads to a complex with a general formula: [Cu(L)2](Y)2. An example of this synthetic strategy is shown in Scheme 4.

Figure 2. Molecular structure of the cation of complex 6: [Cu(dmp)2(CH3CN)]2+. Ellipsoids shown at 35% probability; hydrogen atoms are omitted for the sake of clarity.

distances between the copper center and the fifth ligand, the average distances between the metal center and the nitrogen atoms of the phenanthroline ligands are 2.067 and 2.060 Å for complexes 6 and 4, respectively. The dihedral angle defined by the phenanthroline planes, excluding the metal center, is 65.2°, significantly larger than the corresponding angle found in the chloride-coordinated complex 4. The structural differences between the synthesized complexes are expressed in their corresponding cyclic voltammograms (Figure 3).

Scheme 4. Synthesis of [Cu(dmp)2](ClO4)2a

a

Reagent and conditions: (iii) 2,9-dimethyl-1,10-phenanthroline (2.2 equiv), ethanol, RT, 1 h. Figure 3. Cyclic voltammograms (CVs) of complexes 2, 4, and 6. The measurements were performed in acetonitrile (CH3CN) solution containing 0.1 M [TBA]PF6. Reference electrode: Ag/AgNO3. All the observed waves were chemically reversible. The potential of the Fc+/ Fc redox couple was used as an internal standard (0.63 V vs NHE).

Interestingly, when complex 5 is carefully dried, it appears as a purple powder, which is different from the commonly reported green color of this compound. Upon exposure to moisture, or dissolution in a coordinating solvent, the resulting product turns green. This spectral shift is indicative of a change in coordination sphere of the metal center which alters the spectral properties by distorting the complex symmetry.29 Furthermore, due to the aforementioned tendency of Cu(II) complexes to extend their coordination sphere, the molecular structure complex 5, as resolved from a single crystal grown from an acetonitrile solution, reveals solvent coordination to the metal center yielding the complex 6: [Cu(dmp)2(CH3CN)](ClO4)2 (Figure 2). The phenanthroline-based complex 6 displays a similar geometry to that previously noted for the dichloride complex 4. Once again, the 5-coordinated copper center is out of the phenanthroline planes, resulting in a distorted trigonal bipyramidal geometry. The selected bond distances and angles for complex 6 are listed in Table 1. The Cu−N5 distance for complex 6 (2.040(6) Å) is typical of a bond length between a Cu(II) center and an N-donating ligand, as can be seen by comparison with the other Cu−N phenanthroline distances. The distance to the acetonitrile ligand is significantly shorter than the Cu−Cl distance in compound 4 (2.316(3) Å) due to the smaller radius of the nitrogen atom. Despite the different

Complexes 2 and 4, which are characterized by the coordination of a chloride to the metal center, display an identical cyclic voltammetry. Considering the presence of a significant amount of supporting electrolyte salt ([TBA]PF6) in the CV experiments, it is reasonable to believe that counterion exchange is significant in the case of complex 2, causing the free chloride ion to have negligible influence on the results. The cyclic voltammograms of the chloride coordinated complexes exhibit an oxidation potential (Eox) at 0.24 V (0.87 V vs NHE) and a reduction potential (Ered) at −0.03 V (0.60 V vs NHE). The nonsymmetric and significant peak-to-peak separation of the forward and reverse scan waves (0.27 V) indicates a large charge-transfer resistance associated with slow electron transfer.32−34 It is reasonable to believe that the distorted geometry, caused by the Cu−Cl interaction, undergoes significant structural rearrangement upon reduction/oxidation. The structural rearrangement and the associated reorganization energy are most likely the source of the observed chargetransfer resistance. Interestingly, the half-wave potentials of D

DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry complexes 2 and 4 (0.73 V vs NHE) are remarkably different to that reported for the [Cu(dmp)2]+/[Cu(dmp)2]2+ redox pair in acetonitrile (0.95−1.0 V vs NHE).25,35 Nevertheless, the cyclic voltammogram of the five-coordinated complex 6 is characterized by oxidation and reduction peaks at 0.42 and 0.33 V, respectively (1.05 and 0.96 V vs NHE), giving a half-wave potential (E1/2) equal to 0.38 V (1.00 V vs NHE). Thus, the half-wave potential of complex 6 is drastically different than those of complexes 2 and 4. The similarity between the halfwave potential of complex 6 and that of [Cu(dmp)2]+/ [Cu(dmp)2]2+ may seem counterintuitive at first sight, since one might expect that the coordinated acetonitrile molecule in complex 6 would affect the energetics for electron transfer in the same way as the coordinated chloride does for complexes 2 and 4. However, as explained by Meyer et al., solvent coordination naturally occurs in solution when [Cu(dmp)2]+ is oxidized to its corresponding Cu(II) complex.35 This process takes place due to the aforementioned tendency of such Cu(II) complexes to prefer a 5-coordinated configuration.36 Since the solvent used in the electrochemical experiments is acetonitrile, it can be easily understood that the acetonitrile-coordinated Cu(II) complex (compound 6) will be formed in situ upon oxidation of [Cu(dmp)2]+. Moreover, solvent molecules, such as acetonitrile, display a dynamic decoordination−coordination behavior: when [Cu(dmp)2(CH3CN)]2+ is reduced, the 18electron coordination complex [Cu(dmp)2]+ is formed. The structural changes leave no room for coordination of a fifth ligand (acetonitrile), which therefore dissociates from the metal center. However, according to the data shown in Figure 3, the reduction of the chloride coordinated complexes 2 and 4 does not lead to [Cu(dmp)2]+, as one might presume.37 On the contrary, the electrochemical reversibility of the measured halfwaves suggests that different Cu(I) species, in which the chloride ion remains coordinated to the metal center, forms upon reduction. Several further attempts to electrochemically reduce complex 4 were made by applying more negative potentials, reducing the scan rate, and performing bulk electrolysis (Supporting Information, Figure S1a,b). Despite our efforts, we did not observe the expected half-wave potentials characteristic of the [Cu(dmp)2(CH3CN)]2+/[Cu(dmp)2]+ pair. This led us to the conclusion that the lack of conversion of complexes 2 and 4 into the expected [Cu(dmp)2]+ (which would have resulted in a voltammogram identical to the one of complex 6) observed electrochemically is attributable neither to unsuitable energetics nor to slow kinetics. However, when the chloride-coordinated complex 4 is treated with the chemical reductant L-ascorbic acid (H2Asc), the product obtained is [Cu(dmp)2]PF6 (Scheme 5). L-Ascorbic acid is a common reductant and has been widely employed for the reduction of copper complexes.38−43

Although L-ascorbic acid is extensively used and is of high biological importance, its reduction mechanisms have not yet been fully elucidated.44,45 Ascorbic acid has traditionally been thought of as a one-electron reductant, but its redox reactions almost always involve the loss of both an electron and a proton via proton-coupled electron transfer (PCET).46 Moreover, its oxidation by metal complexes via inner-sphere electron transfer has previously been reported.47−50 In contrast to the electrochemical experiments, chemical reduction by L-ascorbic acid indeed reduced the Cu(II) center in compound 4 to Cu(I) without retaining the chloride ligand. While we find a more detailed investigation of the reduction mechanism fascinating, it is beyond the scope of this work and will be further explored in a separate study. Nevertheless, it is reasonable to assume that the reduction taking place with ascorbic acid has a different, and possibly an inner-sphere, mechanism than that obtained electrochemically.51,52 The crystal structure of complex 7, synthesized according to the direct route shown in Scheme 5, is reported in Figure 4.

Figure 4. Molecular structure of the cation of complex 7 (synthesized according to the route shown in Scheme 5): [Cu(dmp)2]+. Ellipsoids shown at 35% probability; hydrogen atoms are omitted for clarity.

The dihedral angle between the planes defined by the 1,10phenanthroline ligands is 79.9°, which is, as expected, considerably larger than the values determined for the Cu(II) complexes 6 and 4. However, despite its larger dihedral angle, complex 7 still displays a noteworthy distortion from the ideal D2d symmetry (dihedral angle of 90°). Such ligand flattening (toward a square-planar geometry) lowers the symmetry from D2d to D2 and confers the diimine-based Cu(I) complexes their so-called pseudo-tetrahedral geometry.42,35 More details about bond distances and angles of complex 7 can be found in Table 1. The Cu−N distances range from 2.032(4) to 2.068(5) Å, with an average distance of 2.053 Å. As in the previous analyses, the angles around the copper center show considerable distortion from tetrahedral geometry. The N−Cu−N intraligand angles are very similar [82.0(2)−82.3(2)°], while the N−Cu−N interligand angles range from 113.2(2)° to 131.0(2)°, in agreement with the data reported by Coppens et al.53,43 The cyclic voltammogram of complex 7 in acetonitrile is, as expected, identical to that one of complex 6 and is reported in the Supporting Information (Figure S1c).

Scheme 5. Synthesis of [Cu(dmp)2]PF6a

4. CONCLUSION In conclusion, for the synthesis of [Cu(dmp)2]2+ complexes starting from Cu(II) salts, it is imperative to use salts with weakly coordinating counterions such as Cu(ClO4)2, Cu-

a

Reagent and conditions: (iv) L-ascorbic acid (excess), acetonitrile, RT, 1 h. E

DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry (NO3)2, or Cu(BF4)2 and carefully avoid the use of CuCl2 and CuBr2. This work demonstrates that, when starting the synthesis of [Cu(dmp)2]2+ from copper(II) chloride, the resulting product is a 5-coordinated [Cu(dmp)2Cl]+ complex instead of the desired product. Exchanging the counterions by metathesis will solely result in the exchange of the noncoordinated chloride counterion. The coordination of a chloride to the metal center significantly changes the electrochemical properties of the complex with respect to [Cu(dmp)2(CH3CN)]2+ or [Cu(dmp)2]2+. Electrochemical reduction and oxidation of the acetonitrile-coordinated complex is reversible and in agreement with the literature. However, when the chloride is coordinated directly to the copper(II) center instead of a solvent molecule, both the reduction and oxidation occur at more negative potentials and a large peak-to-peak separation is observed. These aspects are of significant importance as, unlike solvent coordination, the chloride ion is not associated with a dynamic decoordinationcoordination behavior upon reduction and oxidation. Furthermore, we observed an interesting effect of the chloridecoordinated complex upon chemical/electrochemical reduction. Electrochemical reduction of [Cu(dmp)2Cl]+ does not yield the expected 4-coordinated [Cu(dmp)2]+ species. In contrast, when L-ascorbic acid was added as a chemical reductant to a solution containing the chloride-coordinated complex, the expected [Cu(dmp)2]+ complex was formed. To the best of our knowledge, despite the vast literature exploring copper complexes prepared from 2,9-disubstituted-1,10-phenanthroline ligands, this is the first time that this behavior has been reported. The present results show that special care must be taken in the chemical characterization of Cu(II) complexes, and most likely also for other multivalent coordination centers, when metal halides are used as initial reagents.

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Lars Kloo: 0000-0002-0168-2942 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The Swedish Research Council and the Swedish Energy Agency are acknowledged for their financial support.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b00225. Detailed investigation of the cyclic voltammetry of complex 4 [Cu(dmp)2Cl]Cl, and complex 7 [Cu(dmp)2]PF6. Apparatus description and reductive bulk electrolysis of complex 4. UV−vis and NIR spectrum of complex 4. UV−vis spectrum of the complex resulting from the reductive bulk electrolysis of complex 4 (PDF) Accession Codes

CCDC 1819088, 1819090, and 1822055 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/ data_request/cif, or by emailing [email protected]. uk, or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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REFERENCES

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*E-mail: [email protected]. ORCID

Hong Chen: 0000-0003-4053-7147 Licheng Sun: 0000-0002-4521-2870 James M. Gardner: 0000-0002-4782-4969 F

DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

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DOI: 10.1021/acs.inorgchem.8b00225 Inorg. Chem. XXXX, XXX, XXX−XXX