Cu(II)-Alkyl Chlorocomplexes: Stable Compounds or Transients? DFT

ARTICLE pubs.acs.org/JPCA

Cu(II)-Alkyl Chlorocomplexes: Stable Compounds or Transients? DFT Prediction of their Structure and EPR Parameters Elena N. Golubeva,*,† Oleg I. Gromov,† and Georgii M. Zhidomirov†,‡ † ‡

Chemistry Department, Lomonosov Moscow State University, Leninskiye Gory 1-3, Moscow, Russia Boreskov Institute of Catalysis (BIC), SO RAN, Novosibirsk, 630090, Russia ABSTRACT: DFT calculations were used for studying the structure and reactivity of organocuprates(II) usually considered as intermediates with very weak Cu
C bond. It was found that calculated principal g-tensor values of model compounds RCu(II)Cl2
are similar to the experimentally found values for organocopper product of photolysis of quaternary ammonium tetrachlorocuprates. The calculations confirm that the most of organocuprates(II) could be stable at ambient conditions, and short lifetimes of organocuprates(II) in solutions or soft matrices are caused by their high reactivity in various bimolecular processes; the rate of those may be close to the rate of diffusion controlled reactions. The charges, spin densities, and d-orbital populations of the Cu atom in them are typical for bivalent copper complexes. Natural bond orbital analysis of organochlorocuprates(II) confirms the formation of polar σ-bond between copper and carbon atoms.

’ INTRODUCTION In 1970s and 1980s, Cu(II) complexes with Cu(II)
C σ-bond were not described in most textbooks and reviews on organocopper compounds.1,2 Sometimes it was even claimed that they did not exist2 in contrast with very rich organometallic chemistry of Cu(I) and Cu(III).1,3
6 Nevertheless, by that time reliable experimental spectral evidence had been obtained showing that organocopper(II) compounds may form under photolysis and radiolysis of Cu(II) complexes, as well as by the reaction of Cu+ ions with organic radicals in water and other solvents, but their lifetimes at ambient conditions usually were very short (10
2
10
6 s),7
9 and they were stable only in frozen solvent matrices at low temperatures. This is why they used to be considered as intermediates with a very weak Cu
C bond.7 Later, a limited group of stable complexes of Cu(II) ions with N-confused porphyrin and its analogues with Cu(II)
C (sp2) bond was synthesized,10
12 and only in the beginning of the twenty-first century a single series of copper(II) halide complexes [CuX(tptm)] (X = F, Cl, Br, I; tptm = tris(2-pyridylthio)methyl) with Cu(II)
C (sp3) bond was obtained and studied by X-ray crystallography, EPR, UV
vis absorption spectroscopy, and DFT calculations.13
15 In both cases, one cannot draw an obvious conclusion about the Cu(II)
C bond strength; in the first case the complexes are immobilized in rigid solvent matrices and in the second one the strong interaction of the central ion with electron donor centers (nitrogen atoms) might play the dominating role in the stabilization. In this paper, we are going to give evidence that short lifetimes of organocuprates(II) in solutions or soft matrices are caused by their high reactivity in various bimolecular processes, and the rate of those may be close to the rate of diffusion controlled reactions.16,17 For this purpose, we carried out quantum-chemical calculations of r 2011 American Chemical Society

structures and principal g-tensor values of model copper(II) complexes with chloride anions and alkyl fragments. The most model compounds were stable at ambient conditions, and theoretical g-tensor values for some structures were close to experimental ones. Calculations also confirmed the primary role of bimolecular mechanisms in vanishing of organocopper(II) compounds. Compounds of general formula RCu(II)Cln1-n or RCu(II)Cln2-n, where R is CH3, iso-C3H7, and (CH3)3N+(CH2)mCHCH3 were chosen as models of organocopper complexes forming under photolysis of tetrachlorocuprates of quaternary ammonium.16

’ CALCULATION DETAILS Unrestricted density functional theory (DFT) calculations of the geometry and electronic structure of model Cu(II)-alkyl chlorocomplexes were performed using the Gaussian 03 program package18 with the PBE exchange-correlation functional.19 Basis set 6-311 g++(3df,3pd) was used for model structures, containing CH3 and iso-C3H7 fragments, and 6-31 g++(d,p) basis set was used for systems, containing (CH3)3N+(CH2)3CHCH3 fragment. The 3-21G basis sets were used for counterions atoms. The geometries of CH3CuCln1-n and [(CH3)3N+(CH2)3CH(CuCl2)
CH3] were optimized both on a 6-311 g++(3df,3pd) level and on a 6-31 g++ (d,p) level. The differences between corresponding bond lengths , and the differences between spin densities did not exceed 0.02 Å were not more than 0.02e (see Tables 1 and 3). The 6-31 g++(d,p) basis was used in the case of more complex models, containing (CH3)3N+(CH2)3CHCH3 units. Single point calculations were Received: March 11, 2011 Revised: June 7, 2011 Published: June 13, 2011 8147

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Table 1. The Cu
Cl and Cu
C Distances Calculated by Unrestricted DFT Approach (Functional PBE)  r, Å complex a

Cu
Cl

Cu
C

CuCl

a

2.06

CuCl2
a

b

2.13

CuCl32-

a

2.33

CH3CuCl

a

2.10

1.92

CH3CuCl2




CH3CuCl32-

Cl
Cu
Cl angle

Cu
Cl
Cl
C dihedral angle

175.8 120

b

2.10

1.92

a

2.20

2.01

144.9

0.2

b

2.21

2.00

143.8

0.2

a

2.40 2.42

2.02 2.00

114.8 114.5

b

C3H7CuCl

a

2.10

1.95

C3H7CuCl2
C3H7CuCl32-

a

2.20

2.06

a

2.40

2.05

(ClC3)Cu

b

2.09

1.95

(Cl2C3)Cu

b

2.18

146.6

0

2.03

150.4

16.3

2.12 2.19

1.97 2.05

147.8

12.5 13.8

2.19 (ClC4)Cu (Cl2C4)Cu

b

(ClC5)Cu

b

2.12

1.96

b

(Cl2C5)Cu


(Cl2C5)Cu + Cl +

NMe4+

(Cl2C5)Cu + Cl


b

2.22; 2.21

2.05

139.5

b

2.22; 2.21

2.05

141.9

3.4

a

2.22

2.07

143.4

9.4

2.06

142.1

9.1

2.19 b

a

2.23; 2.20

6-311++(3df,3pd). b 6-31++(d,p).

performed using energy change convergence criterion of less than 10
8 hartree. The spin unrestricted method was applied even to singlet states when the reaction species are reasonably considered to have an open-shell-singlet electronic configuration. The geometries of model complexes were completely or partially optimized (the distances between methyl groups carbon atoms were fixed in every point of the reaction path scans). Location of minima on the potential energy surface (PES) was checked by calculation of normal-mode frequencies. The stability tests were performed in all cases. The reaction energies were calculated as the difference between the sums of total energies of final and initial states, introducing the zero-point energy corrections. Computed S2 values suggested that spin contamination did not exceed 1
5% except in open-shell-singlet calculations. The scan of potential energy surfaces of the systems, including units with unpaired electrons, was performed using the broken symmetry approach. In this case, S2 quantitatively differed from zero only when distances were more than 3 nm. At shorter lengths, open-shell singlets transformed to closed-shell singlets with S2 values strictly coinciding with theoretical ones. Interpretation of features of electronic structure separately for R and β spin density matrices and calculation of atomic charges was done using the Weinhold natural bond orbital (NBO)20
22 analysis. The NBO analysis gives a better description of the electron distribution in compounds of ionic character, such as those containing metal atoms. It is sensitive for calculations of localized weak interactions, such as charge transfer and hydrogen bonding. The Z-axis in NBO calculations was directed along Cu
C bond. The ORCA program package23 was used to calculate the g-tensors of Cu(II)-organochlorocomplexes. Density functional

theory provides a possibility of calculation of Δg values through the linear response theory (LRT). Calculated in this way Δg values of transition metal complexes can vary depending on exchange-correlation functional and amount of Hartree
Fock exact exchange admixture.24 B(38HF)P86 functional calibrated by Solomon et.al.25 was used for the calculation of Δg values of model Cu(II)-alkyl chlorocomplexes. Δg values of CuCl42- calculated26 using B(38HF)P86 functional are comparable with those obtained with sum-overstates-based approach based on CASPT2 method27 but require less computational effort. In general, a large admixture of exact exchange can make the calculated Δg values of different copper complexes24,26,28,29 close to experimental ones. As it was recommended,30 the CP basis set with three polarization functions31 was used for copper atom, IGLOIII32 was used for chlorine atoms, and IGLO-II31 was used for the rest of the atoms. Calculations of the Δg values utilized the complete mean-field spin
orbit (SO) operator for the Coulomb term.

’ RESULTS AND DISCUSSION Complexes of copper(I) chlorides with ethanolic radicals form as a result of photolysis of tetrachlorocuprate anions [CuCl4]2
by visible light in frozen ethanol matrix.9 The same products were obtained by direct interaction of copper(I) chloride complexes with ethanolic radicals independently generated in a frozen solution by γ-irradiation. Similar complexes also can form by the photolysis of systems [CuCl4]2
-dmf33 and [(CnH2n
1)N+]2[CuCl42
].34 These complexes were assumed by Plyusnin, Bazhin, and Kiseleva9 to be “adducts of Cu(I) with organic radicals”, but they possess 8148

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Figure 1. The geometries of model chloroorganocuprates, PBE/6-311 g++(3df,3pd).

spectral features typical for Cu(II) compounds. For example, the electronic absorption spectra of products of photolysis of [(C6H11)4N+]2[CuCl42
] in frozen solutions by the light with wavelength close to maximum of Cu2+fCl
transition (405 nm) show new bands at 430
450 nm; those can be assigned to nfd (Cl
fCu2+) transitions.9,16 The corresponding EPR spectra also show new signals, belonging to two new Cu(II) complexes.16,17,35 One of the complexes vanishes at 100 K, and its spectrum is not interpreted up to now. The principal values of g-tensor and hyperfine coupling (HFC) tensor A(Cu) of another product of [(C6H11)4N+]2[CuCl42
] photolysis are as follows: g1 = 2.082 ( 0.003, g2 = 2.033 ( 0.001, g3 = 2.024 ( 0.001, A(63Cu)1 = 4.1 ( 0.2 mT, A(63Cu)2 = 3.8 ( 0.1 mT, and A(63Cu)3 = 11.0 ( 0.2 mT. The corresponding halfwidths are ΔH1 = 10.0 mT, ΔH2 = 3.7 mT, and ΔH3 = 6.6 mT. These values are typical for nonsymmetrical Cu(II) compounds.36 It was proposed16,34 that novel cupric complexes contain organic fragments, formed from alkyl substitutes of the quaternary ammonium counterions, because traces of the same free radicals (mainly CH3CH( 3 )CH2
) were registered in the system by EPR, but this assumption was not proved strictly. These compounds disappeared (paramagnetic and colored products vanished) upon heating to 110
115 K. The structure of copper-containing products and mechanism of their vanishing stayed uncertain. Numerous examples of quantum-chemical modeling of Cu(I) (d10) complexes with chloride ions and organic ligands4,37,38

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provided an accurate description of their geometry, electronic structure, and spectral properties. The electronic structure of Cu(II) (d9) chloride complexes with open electron shell was calculated by the use of extended Huckel method,39 density functional theory (DFT),14,25,40,41 and multiconfigurational (CASSCF, CASPT2) methods.27 All these tools provided earlier only qualitative description of the Cu(II) complex geometry, including Jahn
Teller distortions and assignments of transitions in electronic spectra, but recent attempts gave results quantitatively comparable with experiment.25,27,41 For example, density functional theory provides a possibility of calculation of Δg values for a wide range of Cu(II) complexes through the linear response theory (LRT).24,26,41 That is why comparing of experimental principal g-tensor values of the photochemical reactions products with the same values of model structures may help to reveal composition and structures of the complexes. In accordance with experimental data, we proposed the structures of hypothetical copper(II) complexes with chloride anions and alkyl fragments of general formula RCu(II)Cln1
n or RCu(II)Cln2
n, where n = 1
3 (like in all known mononuclear complexes of Cu(I)), and R is CH3, iso-C3H7, and (CH3)3N+ (CH2)mCHCH3, m = 1
3. Complexes with alkyl substitutes are the simplest models, the latter fragments simulate alkyl substitutes formed by quaternary ammonium cations as a result of photolysis of [(CnH2n
1)N+]2[CuCl42
], and the complexes containing them will be denoted below as (ClnCm+2)Cu. The calculated structures of the alkylcopper chlorides displaying bond lengths and bond angles are shown in Figure 1. The Cu
C distances in all species (1.92
2.06 Å) (Table 1) are close to the same values in the stable compounds with Cu(II)
C (sp3) bond, CuX(tptm),14 and are comparable with typical valence distances. Therefore, it is one evidence of a chemical bond formation between copper and carbon atoms. The Cu
C bond elongates with the increase of the number of chlorine atoms in the complex. The Cu
C distance in the complexes with methyl group is shorter than in the complexes with isopropyl group. The Cu
Cl bond length in organocopper compounds is 2
3% longer than Cu
Cl bond length in corresponding Cu(I)Cln1
n anions. The linear complex CuCl2
is twisted upon complexation with radicals; Cl
Cu
Cl angle is 144.9 in CH3CuCl2
and 146.6 in C3H7CuCl2
. Plain Cu(I)Cl32
unit in organocomplexes transforms to pyramidal geometry. Complexes of copper chlorides with [(CH3)3N+(CH2)m CHCH3] are closer similarities of the complexes formed under the photolysis of tetrachlorocuprates of quaternary ammonium.16,17,35 They have some peculiarities comparing with alkyl radical adducts due to positive charge localized on nitrogen atom; certain compounds of formulas RCu(II)Cln2
n, where R = [(CH3)3N+(CH2)mCHCH3], may have cyclic geometry or be unstable. DFT calculations confirm that complexes with three chlorine atoms [(CH3)3N+(CH2)mCH(CuCl3)2
CH3] (Cl3Cm+2) and [(CH3)3N+(CH2)mCH(CuCl3)2
CH3](CH3)4N+ (Cl3Cm+2 + NMe4+) are really unstable and transform to [(CH3)3N+(CH2)m CH(CuCl2)
CH3]Cl
(Cl2Cm+2 + Cl
) and [(CH3)3N+(CH2)m CH(CuCl2)
CH3](CH3)4N+Cl
(Cl2Cm+2 + Cl
+ NMe4+) respectively with the loss of one chloride anion in copper coordination sphere. The structures of (Cl2C5)Cu + Cl
and (Cl2C5)Cu + Cl
+ NMe4+ are presented in Figure 2. Cu
C and Cu
Cl bond lengths in (ClnCm+2)Cu complexes are close to the corresponding values in complexes containing CH3 and isoC3H7 groups. Further, we calculated principal g-tensor values of model compounds and compared them with experimentally acquired g-tensor 8149

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Figure 2. The geometries of chloroorganocuprates [(CH3)3N+(CH2)mCH(Cu(II)Cln)CH3]2
n, n = 1, 2, m = 1, 3, PBE/6-31 g++(d,p).

values of product of photolysis of [(CnH2n-1)N+]2[CuCl4]2
. The Δg values for all methylcopper chlorides (CH3Cu(II)Cln2
n) calculated using DFT approach with B(38HF)P86 hybrid functional significantly differ from those obtained by EPR (Table 2). Complexes with isopropyl group have Δg values reduced with respect to ones in methyl complexes, and in the case of C3H7CuCl2
Δg values are close to the experimentally found ones. Calculated Δg values of the complexes with [(CH3)3N+(CH2)m CHCH3] are strongly affected by chain length m. The m = 3 turned out to be sufficient to get Δg values of model compound containing two chlorine atoms close to experimental Δg values. Introduction of the additional chloride ion into outer sphere of copper changes Δg value not more than by 10%, and this change is close to the calculation error. We suppose, that the EPR spectrum of one of the products of [(C6H11)4N+]2[CuCl42
] photolysis in frozen solutions by the light with wavelength close to maximum of Cu2+fCl
transition (405 nm) obtained by Golubeva et al.16 does belong to organocopper(II) complex. This complex contains two chlorine atoms, because calculated g-tensor values of both model complexes C3H7CuCl2
and (Cl2C5)Cu are similar to experimental ones. That is why the proposition about formation of alkylcopper(II) chlorides as a result of photolysis of tetrachlorocuprates of quaternary ammonium seems plausible. The most stable compounds contain two chlorine atoms, whereas another complex vanishing at lower temperatures may contain distinct number of chlorine atoms in coordination sphere of the central atom.

Table 2. The Principal g Values of the Model Compounds, (DFT/B(38HF)P86) complex

g1

g2

g3

CuClCH3

2.021

2.128

2.188

[CuCl2CH3]


2.003

2.176

2.177

[CuCl3CH3]2


2.037

2.045

2.143

CuClC3H7

2.004

2.037

2.053

[CuCl2C3H7]


2.023

2.032

2.098

[CuCl3C3H7]2


2.003

2.133

2.138

(ClC3)Cu

2.014

2.069

2.109

(Cl2C3)Cu (Cl2C3)Cu+Cl


2.029 2.031

2.037 2.042

2.122 2.129

(ClC4)Cu

2.008

2.038

2.072

(Cl2C4)Cu

2.027

2.035

2.113

(Cl2C4)Cu+Cl


2.029

2.035

2.116

(ClC5)Cu

2.004

2.027

2.044

(Cl2C5)Cu

2.021

2.029

2.088

(Cl2C5)Cu+Cl


2.018

2.027

2.080

Exp[16]

2.024

2.033

2.082

The mechanism of organocuprates(II) vanishing stays obscure. Two alternative mechanisms of organocuprates transformations, mononuclear and binuclear, were proposed.34 The first one relies on the low strength of Cu
C bond and adds up to 8150

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Figure 4. The principal points on the reaction path of CH3CuCl2
with CH3 3 (a) and on the reaction path of two molecules of [N(CH3)4]+2 [CH3Cu(II)Cl2]
(b).

Figure 3. C3H7CuCl2
PES scan along Cu
C distance (a); (Cl2C3) Cu PES scan along Cu
C distance (b).

formation of Cu(I) chlorides and free organic radicals R 3 . Radicals rapidly recombine to stable diamagnetic products RCuðIIÞCln 2
n T CuðIÞCln 2
n + R 3

ðIÞ

This way may be realized at low temperatures (110 K in experiment) and only subject to very low Cu
C bond energy in the complexes (not more than 5
10 kJ/mol). The second mechanism implies bimolecular reactions of organochlorocuprate molecule with itself or with alkyl radicals RCuðIIÞCln 2
n + R 3 f CuðIÞCln 2
n + R 2

ðIIÞ

2RCuðIIÞCln 2
n f 2CuðIÞCln 2
n + R 2

ðIIIÞ

The enthalpies and activation energies of the steps II and III must be negative and negligible, respectively, to allow the reactions to take place at low temperatures. The scans along Cu
C distance in C3H7CuCl2
and in (Cl2C3)Cu are given in Figure 3a,b. The Cu
C bond energy equals 36 and 40 kJ/mol for C3H7CuCl2
and (Cl2C3)Cu, respectively, and the equilibrium I should be significantly shifted to formation of organochlorocuprates even at ambient conditions. That is why we suppose that mononuclear Cu
C bond cleavage is not realized in the reaction conditions. Activation barrier of reaction I for both complexes was not fixed by DFT calculations. As follows from the calculations of the PES of the bimolecular reactions, the interaction of CH3Cu(II)Cl2
with methyl radicals (process II) results in formation of weak complexes of Cu(I)Cl2
with ethane. No activation barrier was observed for this reaction (Figure 4a). The process is exothermic (ΔE = 330 kJ/mol). The

Figure 5. Spin density distribution in (Cl2C5)Cu model structure. Hydrogen atoms are omitted. Notations: Cu, red; Cl, green; C, gray; and N, blue.

same result was obtained for the reaction between two molecules of [N(CH3)4]+2[CH3Cu(II)Cl2]
; the complex [N(CH3)4]+ [Cu(I)Cl2]
...C2H6...[N(CH3)4]+[Cu(I)Cl2]
is generated with energy release (193 kJ/mol) and zero activation energy (Figure 4b). Thereby, the hypothesis about the bimolecular mechanism of chloroorganocuprates(II) vanishing seems to be the most verisimilar and agrees with experimental fact of the disappearance of the EPR signal of the complexes under unfreezing the mobility in 2-chlorobutane solution (∼110 K). Therefore, the spatial isolation of organocuprates(II) might give the opportunity to stabilize them at ambient conditions. 8151

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Table 3. Calculated Natural Atomic Charges and Spin Densities of Model Compounds q (NBO) complex

Cu

CnH2n+1

Cl

Cu

CnH2n+1

Cl

CuCl

a

0.67


0.67

0

CuCl2
a

b

0.58


0.79

0

C3H7CuCl2


a

0.87


0.25


0.76

0.34

0.58

(Cl2C5)Cu

b

0.81

0.61


0.71

0.37

0.46

0.10 (07)

(Cl2C5)Cu+Cl
+NMe4+

b

0.37

0.47

0.08

(Cl2C5)Cu+Cl


b

0.36

0.48

0.08

a

a

spin density (Mulliken)

0.83


0.74

0.54

0 0 0.07 (8)

b

6-311++(3df,3pd). 6-31++(d,p).

Table 4. The NBO description of the bonds between copper and carbon, their populations, and the composition of hybrid orbitals in r- and β-subsystems of RCuCl2
, (R = CH3, C3H7, (CH3)3N+(CH2)3CH( 3 )CH3) complex C3H7CuCl2


(Cl2C5)Cu

R (population)

β (population)

0.28 h(Cu) + 0.96(C) (0.945)

0.83 h(Cu) + 0.57 h(C) (0.982)

h(Cu) = s =
0.98(4s)
0.15(5pz)
0.13(3dx2-y2)

h(Cu) = sd14.25 = 0.26(4s)
0.35(3dx2
y2) + 0.90(3dz2)

h(C) = sp7.29 =
0.34(2s) + 0.93(2pz)

h(C) = sp8.81 =
0.32(2s)
0.94(2pz)

0.32 h(Cu) + 0.95 h(C) (0.959) h(Cu(1)) = s =
0.97(4s)

0.78 h(Cu) + 0.62 h(C) (0.909) h(Cu) = sd1.67 = 0.61(4s)
0.23(3dx2-y2) + 0.75(3dz2)

h(C(2)) = sp5.88 =
0.38(2s)
0.11(2px) + 0.91(2pz)

h(C) = sp8.47 = 0.32(2s) + 0.11(2px)
0.94(2pz)

density in (Cl2C5)Cu is noticeably delocalized between copper, chlorine atoms, and the carbon atom connected with copper atom (Figure 5). According to the NBO analysis, there is a notable charge transfer of about 0.15
0.43 natural atomic charge from the copper chlorides to alkyl fragments. Charges on chlorine atoms do not significantly change (excepting CH3CuCl) (Table 3). Positive charge on copper atom in organocuprates vastly increases by 0.12
0.33e comparing with corresponding Cu(I) complexes, the most significant growth is observed for compounds including CuCl2
and CuCl32
units. Nevertheless, the charge on Cu atom is significantly lower than +2e; this may point out a substantial covalence of metal
ligand bond. Natural electron configurations of Cu ion in chlorocuprates and in corresponding organic complexes are presented below: CuCl CuCl2
C3H7CuCl2
(Cl2C5)Cu

Figure 6. Cu
C natural bonding orbitals in (Cl2C5)Cu. Hydrogen atoms are omitted. Notations: Cu, red; Cl, green; C, gray; and N, blue.

Because of their high reactivity, such compounds may play significant role in radical reactions catalyzed by copper complexes; those are often characterized by the high selectivity.42,43 The knowledge about the electron structure of organocuprates is important for prediction of their chemical properties. That is why we performed the calculations of the spin density and charge distributions of some model organocopper complexes comparing with the corresponding Cu(I) chlorides on the base of the density functional theory and NBO analysis. DFT calculations demonstrate the decrease of spin density on alkyl fragment, the emergence of spin density on the copper atom (up to 0.5e) and on chlorine atoms in organocopper compounds in comparison with Cu(I)Cln1
n (Table 3). The least changes are found for complexes with CuCl2
unit. For example, the spin

[core] 3d9.894s0.404p0.03 [core] 3d9.834s0.554p0.025s0.01 [core] 3d9.644s0.454p0.03 [core] 3d9.654s0.494p0.03

NBO calculations demonstrate that occupancies of 3d orbitals in organocuprates are significantly lower as compared with corresponding Cu(I) chlorides confirming the increase of copper oxidation state. The contributions of 4p orbitals do not depend noticeably on the composition of organocuprates; 4s population decreases, and 3d remains nearly constant with the increase of the number of chloride ions in coordination sphere of copper. The populations of 3d orbital in isopropyl complexes are slightly lower than those in methyl ones. Carbon atom, connected with copper, actually has sp3 hybridization; it possesses with the least p-character (sp3.29) in C3H7CuCl CH3CuCl2
and with the most p-character (sp3.69) in CH3CuCl2
and CH3CuCl32
complexes. As follows from NBO calculations, both R and β electrons participate in the formation of Cu
C covalent bonds in the model complexes containing one chlorine atom and methyl 8152

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The Journal of Physical Chemistry A group. These bonds are virtually polar in R-subsystem, and in the complexes with two or three chloride anions covalent bonds convert into Lewis donor
acceptor interactions. The complexes with two chlorine atoms characterizing with the principle gtensor values similar to experimental ones have polar Cu
C σbonds (Table 4). R-hybrid orbital of copper mainly has s-character, and β-orbital is a combination of 4s, 3dx2
y2, and mostly contributing 3dz2orbital. The visualization of natural bonding orbitals of (Cl2C5)Cu in R and β electronic subsystems clearly displays the σ-character of Cu
C bond (Figure 6).

’ CONCLUSIONS The formation of complexes with the Cu(II)
C σ-bond as a result of photolysis of tetrachlorocuprates of quaternary ammonium was confirmed by the similarity of experimental principle gtensor values and the corresponding values calculated by density functional theory approach for model compounds. As follows from quantum chemical calculations, the most stable Cu(II)
organic compound has two chlorine atoms in the coordination sphere of copper. The less stable complexes may differ from it by quantity of chlorine atoms. The Cu
C bonds are rather strong
ΔECu
C is about 40 kJ/mol in complexes with two chlorine atoms and is even greater in compounds with CuCl and CuCl3 fragments. The interaction of chloroorganocuprates with alkyl radicals results in formation of alkanes and Cu(I) chlorides (reaction II), and no activation barriers were observed for this reaction, as well as for the reaction between two complex molecules (reaction III). These reactions are highly exothermic. High reactivity of chloroorganocuprates may be attributed to the partially covalent character of Cu
C bond. The significant part of spin density localized on alkyl fragment leads to its high reactivity in processes typical for free radicals and to decrease of g-tensor values comparing with tetrachlorocuprates and other ionic Cu complexes. Thus, quantum-chemical calculations corroborate the hypothesis about the bimolecular mechanism of organochlorocuprates(II) decay. The theoretic simulation of the electronic structure of the alkylcopper chlorides indicates that charges, spin densities, d-orbital populations on the Cu ion in them are substantially greater, than corresponding values in Cu(I) chlorocomplexes. These values are typical for bivalent copper complexes. Natural bond orbital analysis of organochlorocuprates(II) confirms the formation of polar σ-bond between copper and carbon atoms. ’ AUTHOR INFORMATION Corresponding Author

*E-mail: (E.N.G.) [email protected].

’ ACKNOWLEDGMENT The research is partially supported by RFBR Grant 10-03-00603a ’ REFERENCES (1) Hathaway, B. J. Copper. In Comprehensive coordination chemistry; Wilkinson, G., Gillard, R. D., McCleverty, J. A., Eds.; Pergamon: Oxford, 1987; Vol. 5, pp 533
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