Use of Crown Ether Functions as Secondary Coordination Spheres for the Manipulation of Ligand–Metal Intramolecular Electron Transfer in Copper–Guanidine Complexes

Abstract Intramolecular electron transfer (IET) between a redox‐active organic ligand and a metal in a complex is of fundamental interest and used in a variety of applications. In this work it is demonstrated that secondary coordination sphere motifs can be applied to trigger a radical change in the electronic structure of copper complexes with a redox‐active guanidine ligand through ligand–metal IET. Hence, crown ether functions attached to the ligand allow the manipulation of the degree of IET between the guanidine ligand and the copper atom through metal encapsulation.


Introduction
Coordination compounds with redox-active organic (or non-innocent) ligandsa re attractive fors everal applications.I nc atalysis, they could act as an electron reservoir that provides electrons for the activation of substrate bonds. [1][2][3][4][5][6][7][8][9][10] Because the different redoxs tates of the ligandsu sually display distinct colors, applications in electrochromic or thermochromic devices could be envisioned. Several applicationsr ely on at unable and reversible intramolecular electron transfer (IET) between the redox-active ligand and the metal. [11] For example, variations in the number of unpaired electrons accompanied by this electron transfer couldb eu sed for the design of switchable magnetic devices, [12] and also allow the tuning of other important materialp roperties such as phase transitions. [13,14] In the past,e xamples for intramolecular ligand-metal electron transfer stimulated by physicalp arameters (temperature, [15][16][17][18] pressure or light irradiation [19][20][21] )w erer eported. [22] Moreover, chemists found ways to trigger ligand-metalI ET by (reversible) chemicalr eactions, for example, at ar emote part of the redoxactive ligand or at the co-ligands attached to the metal. [23] If two or more redox isomersd iffering in their charged istribution are in equilibrium,t he term valence tautomerism is used. [24] The prediction of the electronic structure and the energy barriers for IET in ac oordination compound with one or more redox-active ligandsi so ften difficult, and depends not only on intrinsic properties, but also on the environment (e.g.,t he solvent or the packing of the molecular units in as olid material). [25,26] Fundamental research, both by experimentalists and by theoreticians, is necessary for the advancementi nt his highly promising field of research.
Over the last years, our group studied IET betweenaredoxactive guanidine ligand andc opperi nm ono-a nd dinuclear copper complexes. [6] The strong s-a nd p-donor character of the guanidino groups [27] lead to ad istortion of tetracoordinated copper from the generally preferred square planar coordination mode in the direction to the tetrahedral coordination mode. [28,29] This distortion leads to ad ecrease of the energy barrierfor electron transfer,which in part arises from the different coordination of Cu II and Cu I atoms. [28] In this respect, guanidine ligands act similarly to the thiolate ligands in blue copper proteins. [30] Detailed analysis showedt hat the metal-ligand electron transfer is highly sensitive to changes in the environment (solvent and temperature) anda lso modifications at the co-ligands (hard co-ligands favor the isomer with Cu II and soft co-ligands that with Cu I ). [31][32][33] Te mperature-dependent redox isomerism (valence tautomerism) was observed for somem ononuclear and dinuclear [32] complexes in some selected solvents. Co-ligands ubstitution reactions could trigger IET. [30] First applications of copper complexes with redox-activeg uanidinel igands for cross-coupling reactions between phenolsw ith a non-complementary relationship emerged. [34] An example for temperature-dependent redox isomerism (valence tautomerism) of am ononuclearc opperc omplexw ith the redox-activeg uanidine ligand L Ac is given in Figure 1a. [35,36] Figure 1b shows pairs of copper complexes with the two redox-active bisguanidine ligands L Ac and L Et . [28] The slightly increasedr edoxp otential of L Ac with respectt oL Et leads to a massive effect on the barrier for ligand-metalI ET in the dicationic complexes.H ence for ad icationic copperb romide complex of L Ac valence tautomerism in CH 2 Cl 2 can be observed, whereas no electron transfer occurs in the equivalent complex with L Et . [28] Herein we reporto nc opper complexes with ar edox-active bisguanidine ligand with attached crown ether function, allowing to control the electronic structure of the metal complexes by encapsulation of am etal. Crown ether functions were used in the past to vary the redox potential of transition metal complexes, for example, of ferrocene, [37] cobalt Schiff base complexes, [38] or iron complexes of pyridine-diimines. [39,40] The example in Figure 2a showsaferrocene with ac rown ether group attachedt oone of the cyclopentadienyl rings. [36] The encapsulation of Na + leads to an anodic shift of~60 mV of the Fe III /Fe II reduction potential. Thes econd example in Figure 2a, ap alladium complex with ac rown ether function attached to ab enzenedithiolate ligand, reveals an anodic shift of 100 mV for the ligand-based oxidation upon Na + binding. [41] The electronic structure of the crown ether incorporated Co(salen) complexs hown on the left side of Figure 2b is only slightly af-fected by metal encapsulation. [37] The effect on the redox potential is therefore predominantly an electrostatic effect rather than an inductive effect.F inally,i ron complexes with pyridinediimine ligands( PDI), showing intriguing electronic structures and redox properties (ligand-based oxidation and metal-based reduction), [42] were modified with pendant crowne ther groups (see Lewis structure on the right side of Figure 2b). [38] The complexation of Na + leads to an anodic shift of the ligand potential, but does not change the electronic structure of the complex (S = 0F e II complexw ith doubly reduced PDI ligand).
In this work we show for the first time that intramolecular ligand-metal electron transfer (IET) could be triggered by alkali and earth alkali metal addition to ar emote crowne ther function attached to ar edox-active guanidine ligand. Hence, we reportt he first example showing ar adicalc hange of the electronic structure by meanso fm etal complexation to as econdary coordinations phere; i.e.,f rom aC u II complex with an eutral ligand unit to the redoxi someric Cu I complexw itharadical monocationic ligand unit.
The redox properties of the new ligand were analyzed by cyclic voltammetry (Figure 4). The curve recorded in CH 3 CN solution clearly shows ar eversible redoxe vent at E 1/2 = À0.40 V (E ox = À0.35 V). Due to its similarity with the cyclic voltammograms of relatedl igands in CH 3 CN (Figure 4b), [28,34,45] it is assigned to the redox pair L 2 + /L 0 .I nC H 2 Cl 2 solution,asplitting of the waves into two components emerged, indicating the presence of two one-electron redox steps at slightly different potentials, E 1/2 = À0.35 V( E ox = À0.29 V) for L · + /L 0 and E 1/2 = À0.26 V( E ox = À0.20 V) for L 2 + /L · + .F or comparison, in the case of the related ligand L Ac ,t he two-electron wave observed in CH 3 CN splits much more clearlyi nto two potentially separated one-electron waves in CH 2 Cl 2 solution. [8] The coincidence of first and second oxidation/reduction waves in CH 3 CN could thus be ascribed in partt oap articularly strong solvents tabilization of the dication by the polar CH 3 CN solvent molecules.
Effecto fi ncorporation of K + + and Ba 2 + + into free ligand L Befores tudying copperc omplexes of L,t he effect of alkali (K + ) or earth alkali (Ba 2 + )e ncapsulationb yt he free ligand was inspected. Addition of as olution of KPF 6 (1.5 equivalents) in MeOH toas olutiono fL in CH 2 Cl 2 ,s tirring this mixture for 18 h at room temperature, and aw orkup procedurea fforded paleyellow complex[ K@L](PF 6 )i n8 8% yield. Crystals were grown by diffusion of Et 2 Oi nto as olutiono fC H 2 Cl 2. Figure 5d isplays the solid-state structure from XRD. The K + ion binds to all six oxygen atomso ft he crown ether function, and in addition to two of the fluorines from the PF 6 À counterion. The presence of    the K + ion only slightly varies the bond lengths within the guanidinog roups (see Table 1).
The barium encapsulated compound [Ba@L](OTf) 2 was obtained in 83 %y ield by reaction of L with barium triflate in CH 2 Cl 2 for 18 ha tr oom temperature. The compound was isolated as pale-yellow powder;u nfortunately crystals suitable for an XRD analysiscould not be grown.
The UV/vis spectra of L,[ K@L](PF 6 )a nd [Ba@L](OTf) 2 ,a ll recorded in CH 2 Cl 2 ,a re quite similar (see Supporting Information, Figure S10). Am arginal bathochromic shifto ft he lowest energy band from 324 nm for L to 326 nm for[ K@L](PF 6 )w as observed, whereas ah ypsochromic shift to 321 nm was found for [Ba@L](OTf) 2 .T he cyclic voltammograms in CH 3 CN solution    Figure 6) show reversible redox events, in line with stable metal encapsulationi nt he crown-etherf unction. For [K@L] + ,a two-electron redox process was measured at E 1/2 = À0.35 V (E ox = À0.30 V). For [Ba@L] 2 + ,aredox process at E 1/2 = À0.25 V (E ox = À0.17 V) can be assigned to one-electrono xidation/reduction(in line with the lower current and also reasonable due to the higher charge). Hence the anodics hifts are DE 1/2 = 50 mV for potassium cation encapsulation and DE 1/2 = 150 mV for barium dication encapsulation.I nd ifferencet op reviously reported N-aryl aza-crown ethers, [46] Wursters crown or ortho-Wurster's crownc ompounds, [47] the redox process is reversible before and after metal encapsulation. The measurementsi ndicated that in CH 3 CN or CH 2 Cl 2 solution the bonds between the six oxygen atoms of the ligand and both the K + and the Ba 2 + ion are preserved;a ne quilibrium between free L and [K@L] + /[Ba@L] 2 + was not observed.
The variationo ft he redox potential upon metal encapsulation could arise from inductive and/ore lectrostatic effects. [37] The similarities between the UV/vis spectra recorded before and after metal encapsulation might argue for the domination of electrostatic effects. For ap ure electrostatic effect between ap oint charge and as econd point, the difference in the electric field potential at this point, DE,c ould be estimated from Equation (1):

Copper complexes
Scheme 2g ives an overview of the synthesis of the copper complexes discussed in this work. As already noticedf or L Ac and its copperc omplexes (depicted in Figure 1), UV/vis spectroscopyi sn ot suitable to discriminate betweent he possible redox-states of the ligand unit in the complexes,s ince the differences are too small. Therefore, we predominantly discuss the EPR spectra, the cyclic voltammogramsa nd the solid-state structures, supplemented by the results of quantum-chemical calculations. Firstw er eactedt he new ligand L with Cu(OAc) 2 . The complex [L{Cu(OAc) 2 }] was obtained as ag rey-blue solid in 65 %i solated yield (Scheme 2a). Crystals were grown by diffusion of n-hexane into aT HF solution.T he solid state structure from XRD is visualized in Figure 7. Both acetate groups are essentially h 1 -coordinated, with Cu-O distances of 1.960(3) (Cu-O7) and 1.931(4) (Cu-O9). As econd oxygen of one of the acetate groups establishes in additionaw eak interaction with the copper atom (2.680(4) for Cu-O8). The second oxygen atom of the other acetate group is furthera way from the coppera tom (3.011(4) for CuÀO10). As expected, the imino N = Cb ond distances of the guanidino groups (N1-C7/N4-C12) are elongatedu pon copperc oordination (from 1.282(2)/ 1.286(2) in L to 1.341(7)/1.335 (7) in [L{Cu(OAc) 2 }]). On the other hand, the structural data do not argue for ligand oxidation by IET to the copper atom.
The EPR spectrum of [L{Cu(OAc) 2 }] in CH 3 CN solution shows as ignal for ac opper-centered radical( Cu II complex), with a g value of 2.120 and an A Cu value of 56 G, being in at ypical region for Cu II -guanidine complexes. [28,34] The UV/vis spectrum in CH 2 Cl 2 solution (Supporting Information, Figure S11) is similar to that of the free ligand L.H ence all data indicate that the electronic structure of the complex [L{Cu(OAc) 2 }] is adequately described as Cu II coordinated to an eutral bisguanidine L,b oth in the solid state and in solution.
In subsequent preparative work, ap otassium ion was encapsulated. The complex [K@L{Cu(OAc) 2 }](PF 6 )w as obtained in 61 %i solated yield by reactiono f[ K@L]w ith Cu(OAc) 2 .C rystals were obtainedb yd iffusion of n-pentane into aC H 2 Cl 2 solution. The solid state structure is shown in Figure 8. Again, both ace-tates are essentially h 1 -coordinated (1.928(2) for CuÀO7 and 1.970(2) for CuÀO9). One additional oxygen atom (O10) interacts with the coppera tom (2.643(2) for CuÀO10), while another oxygen (O8) is further away (2.978(2) for CuÀO8) and rather interacts with the potassium atom of an adjacent complex unit (Figure 8b). The CuÀNb ond lengths andt he bond lengths within the guanidino groups are similart ot hose in [L{Cu(OAc) 2 }].T he potassium ion is bound to all six oxygen atoms of the crown ether function, and interactswith two fluorine atoms of the PF 6 À counterion ( Figure 8a). As illustrated in Figure 8b,t he complex units interacti nt he solid state via K···O bonds (2.797(2) )b etween the encapsulatedp otassium atom and one of the acetate ligands of another complex unit, leading to apolymeric structure with zig-zagc hains.
The UV/vis and EPR spectra (Figure 7b)o f[ L{Cu(OAc) 2 }] and [K@L{Cu(OAc) 2 }](PF 6 )a lso look similar. Hencet he data confirm the expectation that the electronic structure of the [L{Cu(OAc) 2 }] complex does not change significantly upon potassium ion encapsulation; aC u II complex with neutral ligand unit prevails.
To favor ligand-metal IET,t he hard acetate co-ligands at the copper atom have to be replaced by softer co-ligands. Therefore, we reacted L with CuCl 2 in CH 3 CN, and obtained the complex [L(CuCl 2 )] in 60 %i solated yield. In this reaction it is important to slowly add CuCl 2 to as olution of L,m aintaining an excesso fL throughout the reaction, since CuCl 2 could coordinate not only to the guanidino groups but also to the oxygen atoms of the crowne ther. [48] Crystals were grown by diffusion of n-hexane into aTHF solution.The solid-state structurei sd isplayed in Figure 9. The bond lengths within the guanidine ligand are in line with the presence of an eutrall igand unit (e.g.,t he elongation of the imino C=Nb onds is not larger than in the coppera cetate complex). Noteworthy,t he dihedral angle at the coppera tom [](CuN 2 ,C uCl 2 )] measures 44.28, being almostp erfectly in-between the angle of 908 for at etrahedral and the angle of 08 for as quare planar coordination mode. The reorganization energy,b eing an important contribution to the barrier for ligand-metalI ET,s hould decrease as a consequence of this special coordinationm ode, [27,28] that arises from the p-donor character of guanidine ligands. [26] The EPR spectrumo f[ L(CuCl 2 )] in the solid state ( Figure S15) is in full agreement with the structure, showing ab road signal at~g = 2.1 due to ac opper-centered radical. Hence, in the solid state the complex is adequately described as aC u II complex with neutralligand unit.
In Figure 10, the EPR spectra of [L(CuCl 2 )] recorded in three solvents, differing in their relative permittivity e r ,a re shown. In   the solvent with the lowest e r value (THF, e r = 7.58), only one broad signal with a g value of 2.112 is found. The shape of the band and its g value clearly argue for ac opper-centeredr adical. Hence,i nT HF the complex [L(CuCl 2 )] is best described as a Cu II complexw ith an eutrall igand L,s imilar to the situation in the solids tate. When the polarity of the solventi si ncreased (CH 2 Cl 2 with al arger e r value of 8.93), as mall signal near g = 2 grows in. In CH 3 CN, as harp signal at g = 2.003 dominates, that is assigned to an organic ligand. Hence in CH 3 CN solution, the complexi sb est described as aC u I complex with an oxidized, radicalm onocationic ligand unit, L · + .T he change in the electronic structure with the solvent polarity could be explained by the charges eparation in the Cu I complex with L · + ligand unit, leadingt oag reater solvent stabilization in polar solvents. Hence, in dependence of the solvent polarity, two different redox isomers (see Lewis structures in Figure 10) are stabilized. In all these experiments, we detected no sign of direct solvent coordination.T he substitution of ac hlorido ligand by acetonitrile or aggregation processes of the mononuclear complex units could be excluded.Q uantum chemicalc alculations (see below) confirm that the change of the relative permittivity is responsible for the changeint he electronic structure.
The cyclic voltammogram of [L(CuCl 2 )] in CH 2 Cl 2 solution (Figure 11 a) shows two reversible redox eventsa tE 1/2 = À0.23 V( E ox = À0.15 V) and E 1/2 = 0.42 V( E ox = 0.50 V), that are assigned to two ligand-centered one-electron redox processes (redox couples L · + /L 0 and L 2 + /L · + ), on the basis of the comparison with the cyclic voltammogram of L and the similarity to [L Ac (CuCl 2 )] (Scheme 1). [34] The reversibility of the redox processes shows that oxidation does not initiate ligand dissociation or aggregation processes via formation of Cu-Cl-Cu bridges.F urthermore, an on-reversible, broad reduction wave (shoulder) at À0.46 Vc an be assigned to copper reduction (redox couple Cu II /Cu I ). Notably,t he first oxidation to [L · + (CuCl 2 )] occurs at ap otential very similar to that of free L.T his is fully consistent with the EPR measurements showing that the ligand unit L in the complex [L(CuCl 2 )] is in its neutral, reduced form in CH 2 Cl 2 solution.
We also recorded CV curves in CH 3 CN solution ( Figure 11 b). Here, as ignificant anodic shift (DE ox = 0.11V)i so bserved for [L(CuCl 2 )] (E ox = À0.24 V) with respectt of ree L (E ox = À0.35 V). The large shift is fully consistent with the formation of the valence tautomer [L · + (CuCl 2 À )] in CH 3 CN solution by ligand-tometal IET,r endering the further oxidation (this time electrochemically) of the ligand unit more difficult. In addition, one recognizes as maller potentiald ifference between the first and second oxidation process in CH 3 CN solutionc ompared with the measurements in CH 2 Cl 2 solution. However,acloser inspection shows that the first oxidation and reduction waves consist of two contributions, which separate more clearly in the cyclic voltammogramso f[ K@L(CuCl 2 )](PF 6 )a nd [Ba@L(CuCl 2 )](OTf) 2 , and mightb ec aused by redox-induced electron transfer (RIET) processes,l eading to the presence of two valence tautomers upon one-electron oxidation (see results obtainedp reviously for [L Ac (CuCl 2 )],S cheme 1, [34] and the discussion below). The reversibility of the redox processes again indicatest hat the solvent CH 3 CN molecules do not initiate changes in the co-ordination mode, for example, by substituting the chlorido ligands.
The complex [K@L(CuCl 2 )](PF 6 )w as prepared in 67 %y ield by reactiono f[ K@L](PF 6 )w ithC uCl 2 .D iffusion of n-pentane into a CH 2 Cl 2 solutiona fforded crystals suitable for as tructural analysis by XRD. Figure 12 illustrates the structure of one [K@L(CuCl 2 )] + unit, and also the interactions between these units through Cu-Cl-K bridgesi nt he solid state. Despite these interactions, there is no significant change in the CuÀCl bond lengths( 2.249(2)/2.219(2) for [L(CuCl 2 )] and 2.263(1)/ 2.232(1) for [K@L(CuCl 2 )](PF 6 )), falling in the typical range of Cu II -bisguanidine complexes. [23] In contrast to [K@L{Cu(OAc 2 )}](PF 6 )t he potassium ion does not interact with the fluorine atoms of the PF 6 À counterion. In the solid state, the cationic units [K@L(CuCl 2 )] + interact with each other by CuÀCl···Ki nteractions, leadingt op olymer chains with K···Cl distanceso f3 .029(1) and 3.399(1) .T he dihedral angle at the copper atom is slightly larger in [K@L(CuCl 2 )] + (53.508)t han in [L(CuCl 2 )] (44.28). However,d ue to the interactions between the [K@L(CuCl 2 )] + units, this angle is expectedt od eviate from the preferredv alue without these interactions (see discussion below). The cyclic voltammogram of [K@L(CuCl 2 )](PF 6 )i nC H 2 Cl 2 is similar to that recorded for [L(CuCl 2 )] (see Figure 11 a). Again, two ligand-centered one-electron redox processes (redox couples L · + /L 0 and L 2 + /L · + )a ppear,a tE 1/2 = À0.21 V( E ox = À0.13 V) and E 1/2 = 0.35 V( E ox = 0.46 V). The non-reversible copper reductiono ccurs at al ower potentialo fÀ0.73 V( redox couple Cu II /Cu I ). The cathodic shift of the Cu II reduction potential could be explained by the reduced electron-donor charactero f L upon K + encapsulation. Hence, the CV curve is in line with an electronic structure with ar educed [K@L] + unit and Cu II . The CV curve recorded for [K@L(CuCl 2 )](PF 6 )i nC H 3 CN solution looks different (see Figure 11 b). In addition to clearly visible oxidation waves at E ox = À0.23 Va nd 0.18 V, as houlder at À0.33 Vappeared. Also, two waves and an additional shoulder showedi nt he direction of reduction. For the complex [L Ac (CuCl 2 )] (see Scheme 1), ar edox-induced electron transfer (RIET) was evidenced. [34] Hence upon one-electrono xidation, IET leads to ac omplex with dicationic ligand unit L Ac 2 + and reduced Cu I metal. Similar RIET processes might lead to am ore complex course of the cyclic voltammetry curve for [K@L(CuCl 2 )](PF 6 )i nC H 3 CN solution.T he reversibility of the redox processes again indicates that oxidation is not accompanied by ac hange of the coordination mode of the copper atom;s ubstitution of chloride by acetonitrile could be excluded on the basis of all experimental results.
In CH 2 Cl 2 solution, the EPR spectra of [L(CuCl 2 )] and [K@L(CuCl 2 )](PF 6 )a re similar, showing ab road signal due to copper-centered radical and only av ery small one due to ligand-centered radical( SupportingI nformation, Figure S14). Hence,t he adequate description is that of aC u II complex with neutrall igand.I nF igure 13, the EPR spectra of [L(CuCl 2 )] and [K@L(CuCl 2 )](PF 6 )i nC H 3 CN solution are compared.B oth spectra display two signals.As harp signal with a g value near 2i s assignedt oar adicalm onocationic ligand unit (L · + ), being much stronger than in CH 2 Cl 2 solution,a nd ab road signal with as ignificantly higher g value assigned to ac opper-centered radical( Cu II ). However,t he ratio of these two signals differs, as estimated by double integration (see Supporting Information, Figure S16). The contribution from the copper-centered radical is much larger for [K@L(CuCl 2 )](PF 6 )t han for [L(CuCl 2 )].T his means that the K + encapsulation by the crown ether function has an effect on the electronic structure of the complex, as it leads to ah igherp reference of the valence tautomer with Cu II and neutrall igand. Moreover,t his is another evidencef or a preserved metal encapsulationi ns olution,b ecause al oss of metal coordination would increase the ionic strength of the solution thereby favoring the other valence tautomer with Cu I and monocationic ligand (see the reportede ffect of salt addition on the valencet automerismo fadinuclearc opperc omplex with redox-active guanidine ligand [32] ).
Finally,w eprepared the complex [Ba@L(CuCl 2 )](OTf) 2 by reacting [Ba@L](OTf) 2 with CuCl 2 in THF solution.T he EPR spectrum of [Ba@L(CuCl 2 )](OTf) 2 in CH 3 CN solution (see Figure13) showst he almoste xclusive presence of copper-centered spin density;o nly av ery weak signal near g = 2( spin density on the organic ligand)i sv isible. This result demonstrates the possibility to massively change the electronic structureo ft he copper  complex by metal coordination at the secondary coordination sphere. To the best of our knowledge,t his is the first authenticated example in which metal encapsulationd oes not only lead to an anodic shift of the ligand and/or metal potential, but to the change of the oxidation state of ar edox-active metal (copper) by IET to the ligand.
The CV curve (recordedi nC H 3 CN solution) displays ac omplicated form with three oxidationw aves. This form might be caused by RIET processes (see also the discussion of the complex with encapsulated K + ion), as provenf or the complex [L Ac (CuCl 2 )],l eading to the presence of a L Ac 2 + ligand unit together with Cu I upon one-electron oxidation, starting with a complex with reduced L Ac ligand and Cu II before oxidation.

DFT calculations
Quantum-chemical calculationsw ere carriedo ut to complement the experimental analysis. Previous calculations with the B3LYP functional and the def2-TZVPb asis set gave reliable results. [28,34] Therefore the calculations in this work also relied on this functional and basis set combination. The solvente ffect was modelled, as in previous work, by the conductor-like screening model (COSMO).
All DFT-calculated structures of [L(CuCl 2 )] and [K@L(CuCl 2 )] + are in good agreementw ith the experimental XRD solid-state structures (see Ta ble S1 in the Supporting Information). The natural bond orbitala nalysis( NBO, Table 2) shows that the spin density for e r = 1i sp redominantly located att he copper atom (ca. 65 %) and the two chlorido ligands(ca. 20 %), leaving only 15 %s pin density for the ligand unit. Hence, the calculated structures for e r = 1c ould safely be described as Cu II complexes with an eutral ligand L,i nl ine with the experimental results of EPR spectroscopy for the solid powder material (see Figure S15 in the Supporting Information).
To investigate the influence of the environment, furtherc alculations were carried out for [L(CuCl 2 )] using COSMO with various relative permittivity valuest om odel the solvent effect, ranging from e r = 1t o4 6.7. At this place we want to stress that one should be cautiousw ith the interpretationo ft he results of these calculations. For some values of the relative permittivity,d ifferent structures were obtained varying only slightly in their energy.M oreover,i ti sn ot clear if we found in all cases the lowest-energy structure. For the plot in Figure14, we consistently used the located structures of lowest energy.F or e r = 1, about 15 %o ft he spin density of [L(CuCl 2 )] resides on the ligand.T his value rises with increasing e r value ( Figure 14 and Ta ble 2). At an e r value of~10, the spin density on the ligand reaches5 0% (dihedrala ngle of~748). For lower e r values, the electronic structure is best described as Cu II with neutral ligand unit (L), and at higher e r values as Cu I with radial monocationic ligand unit (L · + ). The plot in Figure 14 also shows that the dihedrala ngle at the copper atom might increase similarly with the e r value, motivating that this angle is as uitable indicator for the electronic structure. At e r = 37.5, already 75 %o ft he spin density is placedo nt he ligand unit, and the dihedral angle is close to 908.I nl ine with conversionf rom Cu II !Cu I with increasing solventp olarity,t he Cu-N andC u-Cl bond distancesi ncrease (see Supporting Information, Table S1). Not only the spin density and the dihedrala ngle, but also characteristicb ond parameters within the ligand unit signal conversion from neutral to radicalm onocationic ligand unit with increasing e r value (see Supporting Information, Ta ble S1). Hence the N1-C1/N4-C2 bond distances (bondsc onnecting the guanidino groups with the C 6 ring) decrease (from 1.404/1.398 at e r = 1t o1 .358 at e r = 37.5), and the imino N=Cb ond distances within the guanidino groups increase( from 1.321/1.319 at e r = 1t o1 .335/1.336 at e r = 37.5). Moreover, the differences in the CÀCbond distances within the C 6 ring increase.
In previous calculations on [L Ac (CuCl 2 )] (Figure 1), two structures were found at e r = 37.5 with similar energy but distinctly different electronic structure (describable as Cu II complexw ith neutrall igand L Ac and as the redox-isomeric Cu I complexw ith oxidizedl igand L Ac · + ). This result argues for an equilibrium between two redox isomers. Nevertheless, in the calculations a  continuous change of the spin density distribution withi ncreasingr elative permittivity value in the direction toward a ligand-based radicalw as found, similarlyt ot he calculations presented in this work. The distinct signals of organic and metal-based spin density in the EPR spectra at first glance argues for an equilibrium between two redox isomers. However,f urther studies are necessary to decide on the question if the spin density gradually changes with the solventp olarity or if an equilibrium exists between two valencet automers. Anyway,t he B3LYP + COSMO results are in good agreement with the experimental data, both with and without inclusion of the solvent effect. Next, calculations were carriedo ut with an encapsulated crown ether unit (with K + or Ba 2 + ). Both the structuresa nd the spin density contributions, respectively show that the electronic structure of the copper complexes (at e r = 37.5) is significantly changed by the coordination of potassium or barium ions (Figure 15 ba nd c). Before metal encapsulation, as ubstantial part of the spin-density is located on the C 6 ring of the ligand (Figure 15 a), in line with the transfer of electron density from the psystem to the coppera tom. After metal encapsulation, the spin density mainly resides on the copper atom, the directlyb ound two chloride and the two imine nitrogen atoms. The dihedrala ngles are changed from almost perfect tetrahedral toward square planar coordination.H ence the calculations confirm the experimental results, that metal encapsulation by the crown ether function leads to dominant contribution of the Cu II speciesw ith an eutral ligand. Also, the electronic structure does not significantly change with the solvent e r value for the complexes with metal-encapsulated crown ether units, in contrast to the highs ensitivity of the electronic structure toward changes in the solventp olarity for the complex with free crown ether function ( Figure 14, blue lines, and Table 2). Hence, K + or Ba 2 + encapsulationc ausesa massive attenuation of the sensitivity of the electronic structure of the complex toward changesi nt he solventp olarity. For e r = 1, only 16 %o ft he spin density is located on the ligand unit in the complex [K@L(CuCl 2 )] + ,r ising to not more than 24 %a te r = 37.5 ( Figure 14 and Ta ble 2). In polar solvents such as CH 3 CN, metal encapsulation leads to am assive change of the electronic structure, from aC u I complexw ith radicalc ationic ligand (L · + )t oaC u II complex with neutral ligand L (see the illustration in Scheme 3).

Conclusions
In this work crown ether functions were attached as secondary coordination spheres to ar edox-active bisguanidine ligand and the effect of metal encapsulation into the crown ether functions on the electronic structure of copper complexes [L(CuX 2 )] (X = acetate or chloride) of this new ligand L studied. The electronic structure of the copper complex [L(CuCl 2 )] before metal encapsulationc hanges with the solvent polarity.I nn onpolar solvents (CH 2 Cl 2 )a nd in the solid state, the complex is best described asaCu II complex with neutral ligand unit. By contrast, in polar solvents the electronic structure drastically changes to the redox isomeric Cu I complex with radical monocationic ligand,d ue to ligand-metal intramolecular electron transfer (IET). After encapsulation of K + or Ba 2 + ions into the crownether function, the redox isomer assigned to aC u II complex with neutrall igand unit prevails independent of the solvent polarity.W er eport here the first example of ad rastic change in the electronic structure (in polar solvents) through ligandmetal IET by metal encapsulationi nto the crown ether function, going far beyondt he typically observed anodic shift in the ligand potential. The resultss how that ligand-metalI ET could be triggered by coordinationa taremote secondary coordination sphere( Scheme 3). The use of secondary coordination sphere motifs to extensively change the electronic structure of ac oordination compound opens up the possibility for as ophisticated control of the properties and chemical reactivity. Scheme3.Illustrationoft he massive change in the electronic structuret riggered by metal encapsulation into the crown-ether function, as observed in polar solvents. Formally,Ba 2 + encapsulationinitiatesIET from the coppera tom to the ligand (Cu I !Cu II ).