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TD-DFT calculations
Chemical Physics 503 (2018) 14–19
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The electronic spectra and the structures of the individual copper(II) chloride and bromide complexes in acetonitrile according to steady-state absorption spectroscopy and DFT/TD-DFT calculations Pavel K. Olshin a, Olesya S. Myasnikova a, Maria V. Kashina a, Artem O. Gorbunov a, Nikita A. Bogachev a, Viktor O. Kompanets b, Sergey V. Chekalin b, Sergey A. Pulkin a, Vladimir A. Kochemirovsky a, Mikhail Yu. Skripkin a, Andrey S. Mereshchenko a,⇑ a b
Saint-Petersburg State University, 7/9 Universitetskaya nab., St. Petersburg 199034, Russia Institute of Spectroscopy, Russian Academy of Sciences, 5 Fizicheskaya Str., 142190 Troitsk, Moscow, Russia
a r t i c l e
i n f o
Article history: Received 30 May 2017 In final form 27 January 2018 Available online 31 January 2018 Keywords: Copper(II) Halide complexes Chlorocomplexes Bromocomplexes Chloride Bromide Stability constants
a b s t r a c t The results of spectrophotometric study and quantum chemical calculations for copper(II) chloro- and bromocomplexes in acetonitrile are reported. Electronic spectra of the individual copper(II) halide complexes were obtained in a wide spectral range 200–2200 nm. Stability constants of the individual copper (II) halide complexes in acetonitrile were calculated: log b1 = 8.5, log b2 = 15.6, log b3 = 22.5, log b4 = 25.7 for [CuCln]2n and log b1 = 17.0, log b2 = 24.6, log b3 = 28.1, log b4 = 30.4 for [CuBrn]2n. Structures of the studied complexes were optimized and electronic spectra were simulated using DFT and TD-DFT methodologies, respectively. According to the calculations, the more is the number of halide ligands the less is coordination number of copper ion. Ó 2018 Elsevier B.V. All rights reserved.
1. Introduction The distribution of chemical species determines such main characteristics of solution as thermodynamics and kinetics of ligand exchange), stability of the complexes and ‘‘solution-solid phase” equilibrium. Also, local geometry of the transition metal ions defines the electronic structure of the transition metal complexes, which is essential for photochemical processes of transition metal complexes. The copper(II) complexes are known to play important role in catalysis, various biological processes [1] and material science [2]. Therefore, to adequately predict the properties of complex solution systems one should carefully studied the form of the copper(II) complexes. There are numerous debates about the coordination number and the local geometry of copper (II) ions in condensed phase. In solution, copper(II) ions were found to demonstrate coordination numbers equal to four [3] (square planar or distorted tetrahedral), five [4–6] (square pyramidal or trigonal bipyramidal) and six [7–9] (tetrahedrally distorted octahedron (tetragonal bipyramidal)). The copper(II) coordination
⇑ Corresponding author. E-mail address: [email protected] (A.S. Mereshchenko). https://doi.org/10.1016/j.chemphys.2018.01.020 0301-0104/Ó 2018 Elsevier B.V. All rights reserved.
number (CN) and the complex geometry are determined by a solvent properties, nature of the counter ion, ionic strength of the solution, etc. The most common way to determine the complex geometry is to record its UV–vis spectra and to compare them with the ones obtained for the well-studied compounds [10–13]. However, the structure of copper(II) compounds, both spatial and electronic, is complex. To determine exact geometry of the complexes in solution one needs to use combinations of such high-level structural methods as XANES, EXAFS and ab initio quantum chemical calculations [7,14]. There is no single opinion concerning the coordination number and geometry of the copper(II) ions in solution. For example, the most studied copper(II) solvato-complex, namely aquacomplex, was reported to be hexacoordinated [15] and pentacoordinated [5,6] existing in the dynamic vibrational equilibrium between tetragonal bipyramidal and trigonal bipyramidal geometries observed in the last case. The different coordination numbers of copper(II) solvato-complexes are also reported for other solvents. For example, copper(II) solvato-complex in dimethyl sulfoxide ((DMSO) was found to have tetracoordinated square-planar [10] or hexacoordinated axially-elongated distorted-octahedral geometry [7–9]. The problem can be attributed to the strong distortion due to the Jahn-Teller effect. By this reason the axial
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solvent molecule are hard to be detected [3] by modern structural methods. Moreover, the equilibrium between tetra- and hexacoordinated solvato-complexes [17] also makes data analysis difficult. In acetonitrile solutions, the copper(II) solvato-complex is thought to exist as a hexacoordinated distorted octahedron [3,4,18]. However, in some works the copper(II) acetonitrile solvato-complex was found to be tetracoordinated flattened tetrahedron [19] and pentacoordinated elongated square pyramid [20]. Copper(II) halides are ones of the simplest copper(II) model compounds. Therefore, thermodynamics and kinetics of their formation, photochemical properties are intensively studied for such complexes. The complexes were investigated in numerous solvents, like water [18], N,N-dimethylformamide (DMFA) [21,22], dimethyl sulfoxide (DMSO) [16], acetonitrile (MeCN) [23], methanol [24,25], and their mixtures [26]. In solution, a copper(II) ion exists as a solvato-complex and solvent molecule can be substituted by chloride and bromide ligands forming mono, di-, triand tetrahalide complexes. The stability constants, coordination numbers, and local geometries of the halide complexes depend significantly on the solvent. The coordination numbers of copper(II) in solvato-complexes, mono, di-, and trihalide complexes are usually equal six or five, but decrease to four for the tetrahalocomplexes in such solvents as water [27,28], DMSO [16], MeCN [29], DMFA [23]. The stability constants of copper(II) halide complexes decrease with increasing the donor number of solvent. For example, the logarithms of overall stability constants for copper(II) tetrachlorocomplexes were found to be less than 10 in solvents with high donor number or high dielectric permittivity [30] like water [31,32,33] methanol [24,25], ethanol [34], propanol [34], DMSO [16]. In solvents with low donor number and dielectric permittivity such as acetonitrile [23] and propylenecarbonate [10], solvents with, values of logb4 for copper(II) tetrachlorocomplexes exceed 25. It demonstrates that in solvents with high donor ability, copper(II) ions preferably form complexes with the solvents rather than with chloride ions and the solution must contain larger excess of Cl to form the [CuCl4]2 complex. The same tendency is observed in mixed solvents, where addition of even small amounts of donor solvent significantly decreases the stability constants of the copper(II) halide complexes. For example, a small (0.025 mol.%) addition of DMSO (donor number (DN) = 29.8) to acetonitrile (DN = 14.1) decreases the stability constants of the copper(II) chloride complexes in three orders [25]. It means that in such mixtures, copper(II) ions are mainly solvated by DMSO molecules. The stability constants of copper(II) bromocomplexes also depend on the solvent properties, but are usually larger than that of chlorocomplexes, because copper(II) ion can be considered as soft acid in terms of the Hard and Soft Acids and Bases concept, and the bromide ion is softer base than the chloride ion. Thus, for the [CuCl4]2 and [CuBr4]2 complexes logb4 are equal to 5.7 and 6.3 in methanol, 8.3 and 10.3 in ethanol, 10.6 and 10.7 in isopropanol, respectively [24,35]. The stability constants of the aforementioned complexes can be also affected by other factors, such as ionic strength of the solutions where complexation is weaker in presence of great excess of LiClO4 and NH4ClO4 compare to the neat solvent due to the formation of the LiHal and NH4Hal ion pairs [21,22]. Spectral properties of the chloride and bromide complexes are similar, but their absorption bands are shifted. Thus, d-d bands of the copper(II) bromocomplexes are usually slightly red-shifted comparing with chlorocomplexes because Cl ion is located before Br ion in spectrochemical series. The Ligand-to-Metal Charge Transfer (LMCT) bands of the copper(II) bromocomplexes are also shifted towards longer wavelength in respect to the chloride complexes due to more covalent character of Cu-Br than Cu-Cl bond [36]. In this work, the structure and spectral properties of the individual copper(II) chloride and bromide complexes in acetonitrile have been studied using steady-state absorption spectroscopy in
wide UV–vis-NIR (240–2200 nm) spectral range in conjunction with the quantum-chemical DFT/TD-DFT calculations. The stability constants of the copper(II) chloride and bromide complexes were determined and the lowest-energy forms of complexes were revealed. 2. Experimental and computational methods Copper(II) perchlorate hexahydrate (98%), tetraethylammonium perchlorate (>98%), tetraethylammonium bromide (>98%) and acetonitrile (>99.5%) were purchased from Sigma–Aldrich. Tetraethylammonium perchlorate and tetraethylammonium bromide were dried in vacuum oven at 110 °C for 10 h. Acetonitrile was purified and dried by the distillation with the calcium hydride. For spectra measurements, a number of solutions were prepared. The concentration of copper ions from copper(II) perchlorate hexahydrate was 20 mM, the concentration of bromide ions from corresponding tetraethylammonium bromide was varied in range 0–200 mM. Tetraethylammonium perchlorate was added to the solutions to keep the ionic strength constant. The absorption spectra were recorded using a spectrophotometer Lambda 1050 (Perkin Elmer) in 0.2, 1, 10 mm quartz cuvettes. Calculations of the stability constants were performed using ReactLab EQUILIBRIA software. The fraction distributions of the copper (II) halocomplexes were calculated using ‘‘Medusa” software based on the values of the stability constants and the initial component concentrations [37]. The spectra and the equilibrium constants of the chloride complexes were taken from our previous works [38,39]. The geometries of the copper(II) complexes were optimized at the DFT level of theory using B3LYP density functional, which was successfully used in previous work [39]. Also, we oprimized the geometries of copper(II) tetrahalocomplexes using Minesota density functionals (M05, M06, M062x) for comparison. 6-31G(d) basis set was implemented for all calculations. Polarizable continuum model (PCM) was used to simulate solvent media, acetonitrile. For the optimized geometries of the Cu(II) species, vertical excitation transition (VET) energies were calculated using TDDFT (B3LYP/6-31G(d)) methodology. 3. Results and discussion Copper(II) ions exist in the acetonitrile solution as solvatocomplexes [Cu(MeCN)6]2+. Addition of halide ions results in ligand substitution reaction and formation of halide complexes, Eq. (1):
½CuII ðMeCNÞ6
2þ
þ X $ ½CuII ðMeCNÞm X n
2n
þ ð6 mÞMeCNðn ¼ 1 4; m ¼ 0 4; m þ n ¼ 4 6; X ¼ Cl; BrÞ;
ð1Þ
where total coordination numbers can vary from four to six. Each individual halide complex is characterized by its own absorption spectrum. In order to obtain the equilibrium constants and the absorption spectra of the individual bromocomplexes in a wide spectral range, we measured the absorption spectra from 240 to 2200 nm of a set of 16 solutions containing Cu(ClO4)2, NEt4Br, and NEt4ClO4. The concentration of Cu(ClO4)2, the source of copper(II) ions, was kept constant (20 mM). The concentration of NEt4Br, as the source of bromide ions, was varied from 0 to 200 mM (Fig. 1, main text, and S1, Supplementary Materials). The total concentration of NEt4Br and NEt4ClO4 was equal to 200 mM to keep constant the ionic strength of the solution. The complexes exhibit strong Ligand-to-Metal Charge Transfer (LMCT) absorption bands in UV–vis region and weak d-d absorption bands in Near-IR region. LMCT absorption bands correspond
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3.1. Electronic spectra of the individual copper(II) halide complexes
Fig. 1. The UV–vis spectra of 20 mM Cu(ClO4)2 - x mM NEt4Br - (200-x) mM NEt4ClO4 solutions in acetonitrile. Absorbance is normalized to the cuvette path length of 1 cm. Concentration of NEt4Br as a source of bromide ions are shown in legends. The Near-IR part of the spectrum is shown in inset.
to the symmetry allowed electronic transition from the ligandlocalized molecular orbitals to the half-occupied d-orbital of the copper ion. The d-d absorption bands are due to symmetry forbidden electronic transition between the d-orbitals of the copper ion. They appear as vibronic transitions, i.e. coupled with the vibrational modes. Position of the LMCT bands is mostly defined by the ligand properties. Red shift of the LMCT bands in the bromide complexes relative to chloride complexes can be attributed to lower electron affinity of the Br than Cl [36]. Unlike the LMCT bands, the position of weak d-d absorption bands depends on the complex geometry and strength of the ligand field. Obtained spectra were analyzed using ReactLab EQUILIBRIA software, where complex formation was fitted according to equilibrium described by the Eq. (2). As a result, the stability constants (Table 1) and the spectra of the individual copper(II) halide complexes (Fig. 3) were obtained.
Cu2þ þ X $ CuXn2n ðn ¼ 1 4; X ¼ Cl; BrÞ
ð2Þ
The overall stability constants are shown in Table 1. Values of the copper(II) chlorides stability constants adopted from our previous work [37,38] are close to those found by Ishiguro and co-workers (logb1= 9.76, logb2= 17.36, logb3= 22.50, and logb4= 25.44) [20]. The obtained data shows that the copper(II) bromide complexes are significantly more stable than the chloride ones. According to the Hard and Soft Acids and Bases concept, copper(II) ion is the soft acid. Chloride and bromide ions have the same charge, but the radius of the chloride-ion is smaller that of bromide-ion. Therefore, chloride-ion has larger charge density and lower polarizability comparing to bromide-ion, which means that bromide ion is softer base comparing with the chloride ion and, therefore, forms more stable complexes. Based on the values of the stability constants, the fractional distributions of the halide complexes were calculated for the studied concentration range (Fig. 2) using the ‘‘Medusa” software [37].
Table 1 Overall stability constants, bn, of the individual copper(II) halide complexes, [CuXn]2n (X = Cl, Br; n = 1–4) in acetonitrile. Species
log b1
log b2
log b3
log b4
[CuCln]2n [CuBrn]2n
8,5 ± 0.1 17,0 ± 0.1
15,6 ± 0.1 24,6 ± 0.1
22,5 ± 0.1 28,1 ± 0.1
25,7 ± 0.1 30,4 ± 0.1
Spectra of the individual copper(II) chloro- and bromocomplexes obtained by deconvolution of the experimental spectra using the ReactLab EQUILIBRIA software are shown in Fig. 3. Cu2+. Copper(II) exists in acetonitrile solution as a solvato-complex. The weak d-d absorption broad band with the extinction coefficient equals to 20 M1 cm1 is centered at 780 nm. The absorption maximum of LMCT band of the Cu2+ complex is located below 220 nm (e220 = 1370 M1 cm1). [CuHal]+. Copper(II) monohalide complexes in acetonitrile have the only one LMCT absorption band located in UV–vis part of the spectra, and one d-d absorption band in near-IR region, the last one being asymmetric with the broad red shoulder. In case of the CuCl+ complex, the LMCT and d-d absorption bands are located at 292 (e292 = 4190 M1 cm1) and 797 nm (e797 = 77 M1 cm1). For copper(II) bromocomplex the maxima of the LMCT band can be found at 270 nm (emax = 4120 M1 cm1) with shoulder at 330 nm. The d-d absorption band of the CuBr+ is peaking at 832 nm (e832 = 25 M1 cm1). CuHal2. The significant differences between the CuCl2 and CuBr2 can be observed in the UV–vis part of the absorption spectra of the dihalcomplexes. The absorption spectrum of copper(II) dichlorocomplex in acetonitrile demonstrate a broad LMCT absorption band centered at 310 nm (e310 = 4000 M1 cm1) with three shoulders at 265, 380, 460 nm. The copper(II) dibromocomplex demonstrate a very strong LMCT absorption band at 258 nm (e258 = 4875 M1 cm1) with a shoulder at 359 nm (e359 990 M1 cm1). This band also has a low-energy low-intense tail up to 650 nm. In spite of such difference of the LMCT transition energies, the NIR parts of the spectra, corresponding to the d-d transitions, are similar for the chloro- and bromocomplexes. The d-d absorption bands are located at 825 (e825 = 41 M1cm1) and 918 nm (e918 = 86 M1cm1) for the CuCl2 and CuBr2 respectively, but the tail of the bromide complex is much longer than the one of the chloride complex. [CuHal3]. The absorption spectra of the trihalocuprates in acetonitrile have a lot of similarities with the dihalocuprates. In case of the trichlorocomplex, three LMCT absorption bands centered at 258 (e258 = 3252 M1 cm1), 310 (e310 = 4330 M1 cm1) and 460 nm (e460 = 1495 M1 cm1) are observed. The tribromocomplex also have three major LMCT absorption bands peaking at 364 (e364 = 2594 M1 cm1), 434 (e434 = 1522 M1 cm1), and 635 nm (e635 = 1048 M1 cm1). The position of the d-d absorption bands is about the same for both individual complexes, CuCl 3 and CuBr 3 . The d-d absorption bands of trihalocomplexes have significantly different shape comparing with the d-d absorption bands of the other complexes, which may indicate difference in the local geometry of the di- and trihalocomplexes. Two d-d absorption maxima are observed near 900 and 1100 nm. The d-d absorption bands for [CuCl3] are located at 921 (e921 = 83 M1 cm1) and 1089 nm (e1089 = 94 M1 cm1). For [CuBr3], the absorption maxima correspond to 903 (e903 = 110 M1cm1) and 1108 nm (e1108 = 103 M1 cm1). [CuHal4]2. The LMCT spectra of the copper (II) tetrahalocomplexes are complex and contain a large number of absorption bands. The shapes of the LMCT bands of the chloro and bromocomplexes are similar, but bathochromic shift of the copper(II) tetrabromocomplex relatively to the tetrachlorocomplex is observed. Thus, the UV–vis part of the spectrum of tetrachlorocomplex consist of the 242 (2402 M1 cm1), 294 (5954 M1 cm1), 408 nm (2674 M1 cm1) absorption bands and a shoulder at 340 nm (670 M1 cm1). The 408 nm band also has low-energy tail up to 500 nm. The LMCT bands of the [CuBr4]2 have an absorption maxima at 613 (537 M1 cm1), 534 (891 M1 cm1), 351 (2589 M1 cm1) and 275 nm (1503 M1 cm1) as well as a shoulder at 425 nm (660 M1 cm1). The absorption spectra (d-d absorption bands)
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Fig. 2. Fractional distribution of the individual copper(II) chloro- (left panel) and bromocomplexes (right panel) in 20 mM Cu(ClO4)2 – 0–200 mM (C2H5)4NCl acetonitrile solutions.
correspond to the transition from the ligand-localized molecular orbitals to the half-occupied highest energy d-d molecular orbital. Assuming that the copper(II)-centered d-d molecular orbitals are about of the same energy for chloro- and bromocomplexes, the energy difference between the highest energy halide-localized molecular orbital and the copper(II) d-d orbital, corresponding to the lowest-energy LMCT transition, is less for the bromocomplexes. Thereby the low-energy edge of the [CuBrn]2n compounds is red-shifted compare to the [CuCln]2n for the mono-, di-, tri- and tetrasubstituted complexes. At the same time, the position of d-d absorption bands are determined mostly by the ligand field splitting which is affected by the ligand field strength and the complex geometry. Bromide and chloride ions have approximately the same ligand field strength whereas acetonitrile has stronger ligand field. Therefore, the spectra of the copper(II) chloro- and bromocomplexes with the same number of the halide ions are quite similar. Meanwhile, the increasing of the halide ligands number in the first coordination sphere leads to the red-shift of the d-d absorption bands. Thus, the d-d absorption bands of copper(II) solvato-complex are the most blue-shifted. 3.2. The structure of copper(II) halide complexes
Fig. 3. The UV–vis spectra of the individual copper(II) chloro- (CuCln2n (n = 0–4), top panel) [38] and bromocomplexes (CuBrn2n (n = 0–4), bottom panel) in acetonitrile solutions.
of the tetrahalocomplexes are almost identical in NIR part. Thus, [CuCl4]2 and [CuBr4]2 complexes have absorption maxima at 1198 (112 M1 cm1) and 1208 (109 M1 cm1) nm and the long low-energy tails up to 2200 nm. The electron affinity of chloride ions is larger than that of the bromide ions, thus, the molecular orbital of the chloride ligands are located lower in energy. The LMCT absorption bands
The structures and coordination numbers of the copper(II) solvato- and halide complexes are controversial. Thus, the copper(II) acetonitrile solvato-complex structure is usually proposed to be a tetragonal bipyramidal (distorted octahedral) [Cu(MeCN)6]2+ [3,4,14,17]. The copper(II) tetrachloride complex is considered as a distorted tetrahedral [CuCl4]2 in the most works [14,18]. The structure of other halide complexes is unclear and usually considered as a mixture of complexes with different geometries and coordination numbers. To clarify the form of copper(II) chloride and bromide complexes in acetonitrile, the quantum-chemical calculations were performed. The structures of the [CuHalnMeCNm]2n complexes (n = 0–4, Hal = Cl, Br) with coordination numbers equals to 4, 5, and 6 (Table 2) were optimized at the B3LYP/6-31G(d) level of theory, the solvent was simulated by Polarizable Continuum Model (PCM), because acetonitrile does not form a hydrogen bond network. The relative energies of optimized structures are given in Table 3. The geometries of the lowest-energy structures are given in Table S2, Supplementary Materials. Our calculations demonstrate that the more is the number of halide ligands the less is coordination number of copper ion. Thus, the copper(II) solvato-complex in acetonitrile is hexacoordinated, whereas coordination number six is extremely non-favorable for
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Table 2 Possible local geometries of the copper(II) ion for different coordination numbers (CN). Coordination number (CN) Local geometry
6 Tetragonal bipyramidal
5 Trigonal bipyramidal
4 Flattened tetrahedral
Square planar
Scheme
Cu
Cu
Cu
Cu
Table 3 Relative energies (kJ mol1) of copper(II) solvato- and halide complexes calculated using B3LYP/6-31G(d) methodology. Solvent (acetonitrile) was simulated by Polarizable Continuum Model methodology (PCM). Isomers are marked as iso. CN
Local geometry
Cu2+
CuCl+
CuCl2
CuCl 3
CuCl2 4
CuBr+
CuBr2
CuBr 3
CuBr2 4
6 5
Tetragonal bipyramidal Trigonal bipyramidal
0 11.9
2.3 0.80 (iso)
4
Flattened tetrahedral
60.8
29.0
– 0 9.0 (iso) 12.1
– 7.0 10.3 (iso) 0
– 0.6 0 (iso) 27.3
– 8.8 18.1 (iso) 0
63.8
42.7
39.1
38.6
41.8
– 5.1 0 (iso) 12.3 16.3 (iso) 43.6 (cis) 28.5 (trans)
– 0 15.8 (iso) 14.0
Square planar
21.9 4.2 0 (iso) 16.6 18.8 (iso) 42.7 (cis) 33.2 (trans)
45.2
45.8
+ 2 2 CuCl complexes. The 3 , CuCl4 , CuBr , CuBr2, CuBr3 , and CuBr4 monochlorocomplex exists as mixture of penta- and hexacoordinated complexes. Coordination number six is not favorable for all bromocomplexes. (MeCN)5-nBrnCu-NCMe (n = 0–4) coordinate is found to be repulsive. The larger ionic radii of bromide ion (1.82 Å) than chloride ion (1.67 Å) probably results in sterical difficulties for sixth ligand to bind with copper(II) ion. Therefore, monobromocomplex exists in a form of pentacoordinated complex in acetonitrile. Our calculations demonstrate that lowest-energy di- and trihalocomplexes have pentacoordinated trigonal bipyramidal coordination both for chloride and bromide ligands. Pentacoordinated complexes exist as mixture of two closed in energy isomers for each complex (Table 3, main text and Table S1 in Supplementary Materials). For CuCl2 and CuBr2 complexes, the lowest4 4 energy forms correspond to the flattened tetrahedral complexes [CuX4]2. Meanwhile, the differences of relative energies between pentacoordinated trigonal bipyramidal [CuMeCNX4]2 complexes and tetracoordinated flattened tetrahedral complexes [CuX4]2 are less than 9 kJ mol1, and therefore, both forms probably coexist in solution. Thus, the copper(II) tetrahalocomplexes in acetonitrile configuration represents the dynamic equilibrium between the flattened tetrahedron reversibly adding the solvent molecule. Also, we have found that the square pyramidal geometry is not favorable for all considered pentacoordinated complexes, geometry optimization started from the square pyramidal geometry always led to the trigonal bipyramidal geometry. This conclusion is also supported by calculations performed using Minnesota density functionals such as M05, M06, and M062x, where the dominating forms of tetracoordinated copper(II) halide complexes are the bipyramidal [CuMeCNX4]2 complexes and tetracoordinated flattened tetrahedral complexes [CuX4]2 (Table S3 in Supplementary Materials). Other points of interest are the copper(II) di- and trihalocomplexes, which demonstrate the only one stable configuration, namely, the five-coordinated trigonal bipyramid. In the trigonal bipyramidal, five d-d orbitals are split into three levels resulting in two electronic transitions according to the calculated Vertical Excitation Transition energies (Table S2 in Supplementary Materials). This effect is especially pronounced in the copper(II) trihalocomplexes as two d-d-absorption band peaking at about 1100 and 900 nm for both CuCl 3 and CuBr3 complexes. The calculated Vertical Excitation Transition energies of the lowest-energy forms are consistent with the experimentally obtained electronic absorption spectra of the individual complexes (Fig. 3 in the main text
and Table S2 in Supplementary Materials) within the accuracy of 0.2 eV.
4. Conclusions Copper(II) chloro- and bromocomplexes were studied by means of UV–vis-NIR spectrophotometry. The UV–vis-IR absorption spectra and the overall stability constant of the individual copper(II) halide complexes were obtained by the deconvolution of the experimental spectra. The copper(II) bromocomplexes have larger stability constants in acetonitrile compare to the chlorocomplexes. We have found that for the chloro- and bromocomplexes with the same number of the halide ligands, the electronic spectra are similar in the NIR part of the spectra associated with the d-d transitions and significantly different in the UV–vis part of the spectra associated with LMCT transitions. Thus, the lowest-energy LMCT transitions are red-shifted for all bromocomplexes compared to the chlorocomplexes. Our DFT and TD-DFT calculations demonstrate that the more is the number of halide ligands the less is coordination number of copper ion. Thus, the copper(II) solvato-complex in acetonitrile has the octahedral configuration. Most of the mono-, di-, and trihalocomplexes have trigonal bipyramidal local geometry, and the tetrahalocomplexes are flattened tetrahedra. Acknowledgments Authors acknowledge Saint-Petersburg State University for the financial support. Optical measurements were performed at the Center for Optical and Laser Materials Research of the SaintPetersburg State University. Calculations were carried out at the Saint-Petersburg State University Computer Center. This work was supported by the RFBR awards (15-03-05139, and 16-3300646). Authors acknowledge Jplus Consulting for the discount on the Reactlab Equilibria software.
Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at https://doi.org/10.1016/j.chemphys.2018.01. 020.
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The electronic spectra and the structures of the individual copper(II) chloride and bromide complexes in acetonitrile according to steady-state absorption spectroscopy and DFT/TD-DFT calculations Pavel K. Olshin a, Olesya S. Myasnikova a, Maria V. Kashina a, Artem O. Gorbunov a, Nikita A. Bogachev a, Viktor O. Kompanets b, Sergey V. Chekalin b, Sergey A. Pulkin a, Vladimir A. Kochemirovsky a, Mikhail Yu. Skripkin a, Andrey S. Mereshchenko a,⇑ a b
Saint-Petersburg State University, 7/9 Universitetskaya nab., St. Petersburg 199034, Russia Institute of Spectroscopy, Russian Academy of Sciences, 5 Fizicheskaya Str., 142190 Troitsk, Moscow, Russia
a r t i c l e
i n f o
Article history: Received 30 May 2017 In final form 27 January 2018 Available online 31 January 2018 Keywords: Copper(II) Halide complexes Chlorocomplexes Bromocomplexes Chloride Bromide Stability constants
a b s t r a c t The results of spectrophotometric study and quantum chemical calculations for copper(II) chloro- and bromocomplexes in acetonitrile are reported. Electronic spectra of the individual copper(II) halide complexes were obtained in a wide spectral range 200–2200 nm. Stability constants of the individual copper (II) halide complexes in acetonitrile were calculated: log b1 = 8.5, log b2 = 15.6, log b3 = 22.5, log b4 = 25.7 for [CuCln]2n and log b1 = 17.0, log b2 = 24.6, log b3 = 28.1, log b4 = 30.4 for [CuBrn]2n. Structures of the studied complexes were optimized and electronic spectra were simulated using DFT and TD-DFT methodologies, respectively. According to the calculations, the more is the number of halide ligands the less is coordination number of copper ion. Ó 2018 Elsevier B.V. All rights reserved.
1. Introduction The distribution of chemical species determines such main characteristics of solution as thermodynamics and kinetics of ligand exchange), stability of the complexes and ‘‘solution-solid phase” equilibrium. Also, local geometry of the transition metal ions defines the electronic structure of the transition metal complexes, which is essential for photochemical processes of transition metal complexes. The copper(II) complexes are known to play important role in catalysis, various biological processes [1] and material science [2]. Therefore, to adequately predict the properties of complex solution systems one should carefully studied the form of the copper(II) complexes. There are numerous debates about the coordination number and the local geometry of copper (II) ions in condensed phase. In solution, copper(II) ions were found to demonstrate coordination numbers equal to four [3] (square planar or distorted tetrahedral), five [4–6] (square pyramidal or trigonal bipyramidal) and six [7–9] (tetrahedrally distorted octahedron (tetragonal bipyramidal)). The copper(II) coordination
⇑ Corresponding author. E-mail address: [email protected] (A.S. Mereshchenko). https://doi.org/10.1016/j.chemphys.2018.01.020 0301-0104/Ó 2018 Elsevier B.V. All rights reserved.
number (CN) and the complex geometry are determined by a solvent properties, nature of the counter ion, ionic strength of the solution, etc. The most common way to determine the complex geometry is to record its UV–vis spectra and to compare them with the ones obtained for the well-studied compounds [10–13]. However, the structure of copper(II) compounds, both spatial and electronic, is complex. To determine exact geometry of the complexes in solution one needs to use combinations of such high-level structural methods as XANES, EXAFS and ab initio quantum chemical calculations [7,14]. There is no single opinion concerning the coordination number and geometry of the copper(II) ions in solution. For example, the most studied copper(II) solvato-complex, namely aquacomplex, was reported to be hexacoordinated [15] and pentacoordinated [5,6] existing in the dynamic vibrational equilibrium between tetragonal bipyramidal and trigonal bipyramidal geometries observed in the last case. The different coordination numbers of copper(II) solvato-complexes are also reported for other solvents. For example, copper(II) solvato-complex in dimethyl sulfoxide ((DMSO) was found to have tetracoordinated square-planar [10] or hexacoordinated axially-elongated distorted-octahedral geometry [7–9]. The problem can be attributed to the strong distortion due to the Jahn-Teller effect. By this reason the axial
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solvent molecule are hard to be detected [3] by modern structural methods. Moreover, the equilibrium between tetra- and hexacoordinated solvato-complexes [17] also makes data analysis difficult. In acetonitrile solutions, the copper(II) solvato-complex is thought to exist as a hexacoordinated distorted octahedron [3,4,18]. However, in some works the copper(II) acetonitrile solvato-complex was found to be tetracoordinated flattened tetrahedron [19] and pentacoordinated elongated square pyramid [20]. Copper(II) halides are ones of the simplest copper(II) model compounds. Therefore, thermodynamics and kinetics of their formation, photochemical properties are intensively studied for such complexes. The complexes were investigated in numerous solvents, like water [18], N,N-dimethylformamide (DMFA) [21,22], dimethyl sulfoxide (DMSO) [16], acetonitrile (MeCN) [23], methanol [24,25], and their mixtures [26]. In solution, a copper(II) ion exists as a solvato-complex and solvent molecule can be substituted by chloride and bromide ligands forming mono, di-, triand tetrahalide complexes. The stability constants, coordination numbers, and local geometries of the halide complexes depend significantly on the solvent. The coordination numbers of copper(II) in solvato-complexes, mono, di-, and trihalide complexes are usually equal six or five, but decrease to four for the tetrahalocomplexes in such solvents as water [27,28], DMSO [16], MeCN [29], DMFA [23]. The stability constants of copper(II) halide complexes decrease with increasing the donor number of solvent. For example, the logarithms of overall stability constants for copper(II) tetrachlorocomplexes were found to be less than 10 in solvents with high donor number or high dielectric permittivity [30] like water [31,32,33] methanol [24,25], ethanol [34], propanol [34], DMSO [16]. In solvents with low donor number and dielectric permittivity such as acetonitrile [23] and propylenecarbonate [10], solvents with, values of logb4 for copper(II) tetrachlorocomplexes exceed 25. It demonstrates that in solvents with high donor ability, copper(II) ions preferably form complexes with the solvents rather than with chloride ions and the solution must contain larger excess of Cl to form the [CuCl4]2 complex. The same tendency is observed in mixed solvents, where addition of even small amounts of donor solvent significantly decreases the stability constants of the copper(II) halide complexes. For example, a small (0.025 mol.%) addition of DMSO (donor number (DN) = 29.8) to acetonitrile (DN = 14.1) decreases the stability constants of the copper(II) chloride complexes in three orders [25]. It means that in such mixtures, copper(II) ions are mainly solvated by DMSO molecules. The stability constants of copper(II) bromocomplexes also depend on the solvent properties, but are usually larger than that of chlorocomplexes, because copper(II) ion can be considered as soft acid in terms of the Hard and Soft Acids and Bases concept, and the bromide ion is softer base than the chloride ion. Thus, for the [CuCl4]2 and [CuBr4]2 complexes logb4 are equal to 5.7 and 6.3 in methanol, 8.3 and 10.3 in ethanol, 10.6 and 10.7 in isopropanol, respectively [24,35]. The stability constants of the aforementioned complexes can be also affected by other factors, such as ionic strength of the solutions where complexation is weaker in presence of great excess of LiClO4 and NH4ClO4 compare to the neat solvent due to the formation of the LiHal and NH4Hal ion pairs [21,22]. Spectral properties of the chloride and bromide complexes are similar, but their absorption bands are shifted. Thus, d-d bands of the copper(II) bromocomplexes are usually slightly red-shifted comparing with chlorocomplexes because Cl ion is located before Br ion in spectrochemical series. The Ligand-to-Metal Charge Transfer (LMCT) bands of the copper(II) bromocomplexes are also shifted towards longer wavelength in respect to the chloride complexes due to more covalent character of Cu-Br than Cu-Cl bond [36]. In this work, the structure and spectral properties of the individual copper(II) chloride and bromide complexes in acetonitrile have been studied using steady-state absorption spectroscopy in
wide UV–vis-NIR (240–2200 nm) spectral range in conjunction with the quantum-chemical DFT/TD-DFT calculations. The stability constants of the copper(II) chloride and bromide complexes were determined and the lowest-energy forms of complexes were revealed. 2. Experimental and computational methods Copper(II) perchlorate hexahydrate (98%), tetraethylammonium perchlorate (>98%), tetraethylammonium bromide (>98%) and acetonitrile (>99.5%) were purchased from Sigma–Aldrich. Tetraethylammonium perchlorate and tetraethylammonium bromide were dried in vacuum oven at 110 °C for 10 h. Acetonitrile was purified and dried by the distillation with the calcium hydride. For spectra measurements, a number of solutions were prepared. The concentration of copper ions from copper(II) perchlorate hexahydrate was 20 mM, the concentration of bromide ions from corresponding tetraethylammonium bromide was varied in range 0–200 mM. Tetraethylammonium perchlorate was added to the solutions to keep the ionic strength constant. The absorption spectra were recorded using a spectrophotometer Lambda 1050 (Perkin Elmer) in 0.2, 1, 10 mm quartz cuvettes. Calculations of the stability constants were performed using ReactLab EQUILIBRIA software. The fraction distributions of the copper (II) halocomplexes were calculated using ‘‘Medusa” software based on the values of the stability constants and the initial component concentrations [37]. The spectra and the equilibrium constants of the chloride complexes were taken from our previous works [38,39]. The geometries of the copper(II) complexes were optimized at the DFT level of theory using B3LYP density functional, which was successfully used in previous work [39]. Also, we oprimized the geometries of copper(II) tetrahalocomplexes using Minesota density functionals (M05, M06, M062x) for comparison. 6-31G(d) basis set was implemented for all calculations. Polarizable continuum model (PCM) was used to simulate solvent media, acetonitrile. For the optimized geometries of the Cu(II) species, vertical excitation transition (VET) energies were calculated using TDDFT (B3LYP/6-31G(d)) methodology. 3. Results and discussion Copper(II) ions exist in the acetonitrile solution as solvatocomplexes [Cu(MeCN)6]2+. Addition of halide ions results in ligand substitution reaction and formation of halide complexes, Eq. (1):
½CuII ðMeCNÞ6
2þ
þ X $ ½CuII ðMeCNÞm X n
2n
þ ð6 mÞMeCNðn ¼ 1 4; m ¼ 0 4; m þ n ¼ 4 6; X ¼ Cl; BrÞ;
ð1Þ
where total coordination numbers can vary from four to six. Each individual halide complex is characterized by its own absorption spectrum. In order to obtain the equilibrium constants and the absorption spectra of the individual bromocomplexes in a wide spectral range, we measured the absorption spectra from 240 to 2200 nm of a set of 16 solutions containing Cu(ClO4)2, NEt4Br, and NEt4ClO4. The concentration of Cu(ClO4)2, the source of copper(II) ions, was kept constant (20 mM). The concentration of NEt4Br, as the source of bromide ions, was varied from 0 to 200 mM (Fig. 1, main text, and S1, Supplementary Materials). The total concentration of NEt4Br and NEt4ClO4 was equal to 200 mM to keep constant the ionic strength of the solution. The complexes exhibit strong Ligand-to-Metal Charge Transfer (LMCT) absorption bands in UV–vis region and weak d-d absorption bands in Near-IR region. LMCT absorption bands correspond
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3.1. Electronic spectra of the individual copper(II) halide complexes
Fig. 1. The UV–vis spectra of 20 mM Cu(ClO4)2 - x mM NEt4Br - (200-x) mM NEt4ClO4 solutions in acetonitrile. Absorbance is normalized to the cuvette path length of 1 cm. Concentration of NEt4Br as a source of bromide ions are shown in legends. The Near-IR part of the spectrum is shown in inset.
to the symmetry allowed electronic transition from the ligandlocalized molecular orbitals to the half-occupied d-orbital of the copper ion. The d-d absorption bands are due to symmetry forbidden electronic transition between the d-orbitals of the copper ion. They appear as vibronic transitions, i.e. coupled with the vibrational modes. Position of the LMCT bands is mostly defined by the ligand properties. Red shift of the LMCT bands in the bromide complexes relative to chloride complexes can be attributed to lower electron affinity of the Br than Cl [36]. Unlike the LMCT bands, the position of weak d-d absorption bands depends on the complex geometry and strength of the ligand field. Obtained spectra were analyzed using ReactLab EQUILIBRIA software, where complex formation was fitted according to equilibrium described by the Eq. (2). As a result, the stability constants (Table 1) and the spectra of the individual copper(II) halide complexes (Fig. 3) were obtained.
Cu2þ þ X $ CuXn2n ðn ¼ 1 4; X ¼ Cl; BrÞ
ð2Þ
The overall stability constants are shown in Table 1. Values of the copper(II) chlorides stability constants adopted from our previous work [37,38] are close to those found by Ishiguro and co-workers (logb1= 9.76, logb2= 17.36, logb3= 22.50, and logb4= 25.44) [20]. The obtained data shows that the copper(II) bromide complexes are significantly more stable than the chloride ones. According to the Hard and Soft Acids and Bases concept, copper(II) ion is the soft acid. Chloride and bromide ions have the same charge, but the radius of the chloride-ion is smaller that of bromide-ion. Therefore, chloride-ion has larger charge density and lower polarizability comparing to bromide-ion, which means that bromide ion is softer base comparing with the chloride ion and, therefore, forms more stable complexes. Based on the values of the stability constants, the fractional distributions of the halide complexes were calculated for the studied concentration range (Fig. 2) using the ‘‘Medusa” software [37].
Table 1 Overall stability constants, bn, of the individual copper(II) halide complexes, [CuXn]2n (X = Cl, Br; n = 1–4) in acetonitrile. Species
log b1
log b2
log b3
log b4
[CuCln]2n [CuBrn]2n
8,5 ± 0.1 17,0 ± 0.1
15,6 ± 0.1 24,6 ± 0.1
22,5 ± 0.1 28,1 ± 0.1
25,7 ± 0.1 30,4 ± 0.1
Spectra of the individual copper(II) chloro- and bromocomplexes obtained by deconvolution of the experimental spectra using the ReactLab EQUILIBRIA software are shown in Fig. 3. Cu2+. Copper(II) exists in acetonitrile solution as a solvato-complex. The weak d-d absorption broad band with the extinction coefficient equals to 20 M1 cm1 is centered at 780 nm. The absorption maximum of LMCT band of the Cu2+ complex is located below 220 nm (e220 = 1370 M1 cm1). [CuHal]+. Copper(II) monohalide complexes in acetonitrile have the only one LMCT absorption band located in UV–vis part of the spectra, and one d-d absorption band in near-IR region, the last one being asymmetric with the broad red shoulder. In case of the CuCl+ complex, the LMCT and d-d absorption bands are located at 292 (e292 = 4190 M1 cm1) and 797 nm (e797 = 77 M1 cm1). For copper(II) bromocomplex the maxima of the LMCT band can be found at 270 nm (emax = 4120 M1 cm1) with shoulder at 330 nm. The d-d absorption band of the CuBr+ is peaking at 832 nm (e832 = 25 M1 cm1). CuHal2. The significant differences between the CuCl2 and CuBr2 can be observed in the UV–vis part of the absorption spectra of the dihalcomplexes. The absorption spectrum of copper(II) dichlorocomplex in acetonitrile demonstrate a broad LMCT absorption band centered at 310 nm (e310 = 4000 M1 cm1) with three shoulders at 265, 380, 460 nm. The copper(II) dibromocomplex demonstrate a very strong LMCT absorption band at 258 nm (e258 = 4875 M1 cm1) with a shoulder at 359 nm (e359 990 M1 cm1). This band also has a low-energy low-intense tail up to 650 nm. In spite of such difference of the LMCT transition energies, the NIR parts of the spectra, corresponding to the d-d transitions, are similar for the chloro- and bromocomplexes. The d-d absorption bands are located at 825 (e825 = 41 M1cm1) and 918 nm (e918 = 86 M1cm1) for the CuCl2 and CuBr2 respectively, but the tail of the bromide complex is much longer than the one of the chloride complex. [CuHal3]. The absorption spectra of the trihalocuprates in acetonitrile have a lot of similarities with the dihalocuprates. In case of the trichlorocomplex, three LMCT absorption bands centered at 258 (e258 = 3252 M1 cm1), 310 (e310 = 4330 M1 cm1) and 460 nm (e460 = 1495 M1 cm1) are observed. The tribromocomplex also have three major LMCT absorption bands peaking at 364 (e364 = 2594 M1 cm1), 434 (e434 = 1522 M1 cm1), and 635 nm (e635 = 1048 M1 cm1). The position of the d-d absorption bands is about the same for both individual complexes, CuCl 3 and CuBr 3 . The d-d absorption bands of trihalocomplexes have significantly different shape comparing with the d-d absorption bands of the other complexes, which may indicate difference in the local geometry of the di- and trihalocomplexes. Two d-d absorption maxima are observed near 900 and 1100 nm. The d-d absorption bands for [CuCl3] are located at 921 (e921 = 83 M1 cm1) and 1089 nm (e1089 = 94 M1 cm1). For [CuBr3], the absorption maxima correspond to 903 (e903 = 110 M1cm1) and 1108 nm (e1108 = 103 M1 cm1). [CuHal4]2. The LMCT spectra of the copper (II) tetrahalocomplexes are complex and contain a large number of absorption bands. The shapes of the LMCT bands of the chloro and bromocomplexes are similar, but bathochromic shift of the copper(II) tetrabromocomplex relatively to the tetrachlorocomplex is observed. Thus, the UV–vis part of the spectrum of tetrachlorocomplex consist of the 242 (2402 M1 cm1), 294 (5954 M1 cm1), 408 nm (2674 M1 cm1) absorption bands and a shoulder at 340 nm (670 M1 cm1). The 408 nm band also has low-energy tail up to 500 nm. The LMCT bands of the [CuBr4]2 have an absorption maxima at 613 (537 M1 cm1), 534 (891 M1 cm1), 351 (2589 M1 cm1) and 275 nm (1503 M1 cm1) as well as a shoulder at 425 nm (660 M1 cm1). The absorption spectra (d-d absorption bands)
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Fig. 2. Fractional distribution of the individual copper(II) chloro- (left panel) and bromocomplexes (right panel) in 20 mM Cu(ClO4)2 – 0–200 mM (C2H5)4NCl acetonitrile solutions.
correspond to the transition from the ligand-localized molecular orbitals to the half-occupied highest energy d-d molecular orbital. Assuming that the copper(II)-centered d-d molecular orbitals are about of the same energy for chloro- and bromocomplexes, the energy difference between the highest energy halide-localized molecular orbital and the copper(II) d-d orbital, corresponding to the lowest-energy LMCT transition, is less for the bromocomplexes. Thereby the low-energy edge of the [CuBrn]2n compounds is red-shifted compare to the [CuCln]2n for the mono-, di-, tri- and tetrasubstituted complexes. At the same time, the position of d-d absorption bands are determined mostly by the ligand field splitting which is affected by the ligand field strength and the complex geometry. Bromide and chloride ions have approximately the same ligand field strength whereas acetonitrile has stronger ligand field. Therefore, the spectra of the copper(II) chloro- and bromocomplexes with the same number of the halide ions are quite similar. Meanwhile, the increasing of the halide ligands number in the first coordination sphere leads to the red-shift of the d-d absorption bands. Thus, the d-d absorption bands of copper(II) solvato-complex are the most blue-shifted. 3.2. The structure of copper(II) halide complexes
Fig. 3. The UV–vis spectra of the individual copper(II) chloro- (CuCln2n (n = 0–4), top panel) [38] and bromocomplexes (CuBrn2n (n = 0–4), bottom panel) in acetonitrile solutions.
of the tetrahalocomplexes are almost identical in NIR part. Thus, [CuCl4]2 and [CuBr4]2 complexes have absorption maxima at 1198 (112 M1 cm1) and 1208 (109 M1 cm1) nm and the long low-energy tails up to 2200 nm. The electron affinity of chloride ions is larger than that of the bromide ions, thus, the molecular orbital of the chloride ligands are located lower in energy. The LMCT absorption bands
The structures and coordination numbers of the copper(II) solvato- and halide complexes are controversial. Thus, the copper(II) acetonitrile solvato-complex structure is usually proposed to be a tetragonal bipyramidal (distorted octahedral) [Cu(MeCN)6]2+ [3,4,14,17]. The copper(II) tetrachloride complex is considered as a distorted tetrahedral [CuCl4]2 in the most works [14,18]. The structure of other halide complexes is unclear and usually considered as a mixture of complexes with different geometries and coordination numbers. To clarify the form of copper(II) chloride and bromide complexes in acetonitrile, the quantum-chemical calculations were performed. The structures of the [CuHalnMeCNm]2n complexes (n = 0–4, Hal = Cl, Br) with coordination numbers equals to 4, 5, and 6 (Table 2) were optimized at the B3LYP/6-31G(d) level of theory, the solvent was simulated by Polarizable Continuum Model (PCM), because acetonitrile does not form a hydrogen bond network. The relative energies of optimized structures are given in Table 3. The geometries of the lowest-energy structures are given in Table S2, Supplementary Materials. Our calculations demonstrate that the more is the number of halide ligands the less is coordination number of copper ion. Thus, the copper(II) solvato-complex in acetonitrile is hexacoordinated, whereas coordination number six is extremely non-favorable for
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P.K. Olshin et al. / Chemical Physics 503 (2018) 14–19
Table 2 Possible local geometries of the copper(II) ion for different coordination numbers (CN). Coordination number (CN) Local geometry
6 Tetragonal bipyramidal
5 Trigonal bipyramidal
4 Flattened tetrahedral
Square planar
Scheme
Cu
Cu
Cu
Cu
Table 3 Relative energies (kJ mol1) of copper(II) solvato- and halide complexes calculated using B3LYP/6-31G(d) methodology. Solvent (acetonitrile) was simulated by Polarizable Continuum Model methodology (PCM). Isomers are marked as iso. CN
Local geometry
Cu2+
CuCl+
CuCl2
CuCl 3
CuCl2 4
CuBr+
CuBr2
CuBr 3
CuBr2 4
6 5
Tetragonal bipyramidal Trigonal bipyramidal
0 11.9
2.3 0.80 (iso)
4
Flattened tetrahedral
60.8
29.0
– 0 9.0 (iso) 12.1
– 7.0 10.3 (iso) 0
– 0.6 0 (iso) 27.3
– 8.8 18.1 (iso) 0
63.8
42.7
39.1
38.6
41.8
– 5.1 0 (iso) 12.3 16.3 (iso) 43.6 (cis) 28.5 (trans)
– 0 15.8 (iso) 14.0
Square planar
21.9 4.2 0 (iso) 16.6 18.8 (iso) 42.7 (cis) 33.2 (trans)
45.2
45.8
+ 2 2 CuCl complexes. The 3 , CuCl4 , CuBr , CuBr2, CuBr3 , and CuBr4 monochlorocomplex exists as mixture of penta- and hexacoordinated complexes. Coordination number six is not favorable for all bromocomplexes. (MeCN)5-nBrnCu-NCMe (n = 0–4) coordinate is found to be repulsive. The larger ionic radii of bromide ion (1.82 Å) than chloride ion (1.67 Å) probably results in sterical difficulties for sixth ligand to bind with copper(II) ion. Therefore, monobromocomplex exists in a form of pentacoordinated complex in acetonitrile. Our calculations demonstrate that lowest-energy di- and trihalocomplexes have pentacoordinated trigonal bipyramidal coordination both for chloride and bromide ligands. Pentacoordinated complexes exist as mixture of two closed in energy isomers for each complex (Table 3, main text and Table S1 in Supplementary Materials). For CuCl2 and CuBr2 complexes, the lowest4 4 energy forms correspond to the flattened tetrahedral complexes [CuX4]2. Meanwhile, the differences of relative energies between pentacoordinated trigonal bipyramidal [CuMeCNX4]2 complexes and tetracoordinated flattened tetrahedral complexes [CuX4]2 are less than 9 kJ mol1, and therefore, both forms probably coexist in solution. Thus, the copper(II) tetrahalocomplexes in acetonitrile configuration represents the dynamic equilibrium between the flattened tetrahedron reversibly adding the solvent molecule. Also, we have found that the square pyramidal geometry is not favorable for all considered pentacoordinated complexes, geometry optimization started from the square pyramidal geometry always led to the trigonal bipyramidal geometry. This conclusion is also supported by calculations performed using Minnesota density functionals such as M05, M06, and M062x, where the dominating forms of tetracoordinated copper(II) halide complexes are the bipyramidal [CuMeCNX4]2 complexes and tetracoordinated flattened tetrahedral complexes [CuX4]2 (Table S3 in Supplementary Materials). Other points of interest are the copper(II) di- and trihalocomplexes, which demonstrate the only one stable configuration, namely, the five-coordinated trigonal bipyramid. In the trigonal bipyramidal, five d-d orbitals are split into three levels resulting in two electronic transitions according to the calculated Vertical Excitation Transition energies (Table S2 in Supplementary Materials). This effect is especially pronounced in the copper(II) trihalocomplexes as two d-d-absorption band peaking at about 1100 and 900 nm for both CuCl 3 and CuBr3 complexes. The calculated Vertical Excitation Transition energies of the lowest-energy forms are consistent with the experimentally obtained electronic absorption spectra of the individual complexes (Fig. 3 in the main text
and Table S2 in Supplementary Materials) within the accuracy of 0.2 eV.
4. Conclusions Copper(II) chloro- and bromocomplexes were studied by means of UV–vis-NIR spectrophotometry. The UV–vis-IR absorption spectra and the overall stability constant of the individual copper(II) halide complexes were obtained by the deconvolution of the experimental spectra. The copper(II) bromocomplexes have larger stability constants in acetonitrile compare to the chlorocomplexes. We have found that for the chloro- and bromocomplexes with the same number of the halide ligands, the electronic spectra are similar in the NIR part of the spectra associated with the d-d transitions and significantly different in the UV–vis part of the spectra associated with LMCT transitions. Thus, the lowest-energy LMCT transitions are red-shifted for all bromocomplexes compared to the chlorocomplexes. Our DFT and TD-DFT calculations demonstrate that the more is the number of halide ligands the less is coordination number of copper ion. Thus, the copper(II) solvato-complex in acetonitrile has the octahedral configuration. Most of the mono-, di-, and trihalocomplexes have trigonal bipyramidal local geometry, and the tetrahalocomplexes are flattened tetrahedra. Acknowledgments Authors acknowledge Saint-Petersburg State University for the financial support. Optical measurements were performed at the Center for Optical and Laser Materials Research of the SaintPetersburg State University. Calculations were carried out at the Saint-Petersburg State University Computer Center. This work was supported by the RFBR awards (15-03-05139, and 16-3300646). Authors acknowledge Jplus Consulting for the discount on the Reactlab Equilibria software.
Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at https://doi.org/10.1016/j.chemphys.2018.01. 020.
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