The Importance of Pi-Interactions in Crystal Engineering (Frontiers in Crystal Engineering) ||...

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JWST146-c10 JWST146-Tiekink January 20, 2012 13:4 Printer: Yet to come

10Noncovalent Interactions of

π -Systems in Crystal Structures ofTransition-Metal Complexes

Goran V. Janjic1 and Snezana D. Zaric2

1Institute of Chemistry, Technology and Metallurgy, Njegoseva 12, University ofBelgrade, Belgrade, Serbia

2Department of Chemistry, University of Belgrade, Studentski trg 12-16,Belgrade, Serbia

10.1 Introduction

In the last decade several noncovalent interactions involving π -systems that are quitespecific to the crystal structures of transition-metal complexes or organometallic compoundshave been unveiled. These noncovalent interactions can involve two types of π -systems.The first type comprises π -systems of organic molecules, often aromatic, whereas thesecond type comprises planar chelate rings with delocalised π -bonds.

The first type of π -systems, organic molecules, interact with ligands coordinated toa metal. The interactions can be classified as metal–ligand aromatic cation-π interac-tions (MLACπ ) and metal-ligand XH. . .π (MLXH. . .π ) interactions. In MLACπ inter-actions, a cationic transition-metal complex interacts with a π -system. Depending on thecharge of the complex, interactions can be very strong, up to 30 kcal/mol. In all describedcases of MLACπ interactions, the hydrogen atoms of the ligand are oriented towards theπ -system. Accordingly, these interactions can be considered also as charge-assisted XH. . .π

interactions.

The Importance of Pi-Interactions in Crystal Engineering: Frontiers in Crystal Engineering, First Edition.Edited by Edward R. T. Tiekink and Julio Zukerman-Schpector.© 2012 John Wiley & Sons, Ltd. Published 2012 by John Wiley & Sons, Ltd.

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244 The Importance of Pi-Interactions in Crystal Engineering

If the transition-metal complex is neutral and interacts with a π -system, the hydrogenatom of a ligand again interacts with a π -system but, these interactions are in fact a subclassof XH. . .π interactions; metal ligand XH. . .π (MLXH. . .π ) interactions. By coordinatingto a metal ion, ligands attain some positive charge with a concomitant increase of posi-tive charge on ligand-bound hydrogen atoms. For example, coordinated water moleculesform stronger XH. . .π interactions than noncoordinated water molecules. Thus, one cananticipate that XH. . .π interactions of ligands can be quite strong.

The second type of π -system that can participate in noncovalent interactions to bediscussed herein are planar chelate rings with delocalized π -bonds. Aromatic or not aro-matic, they have π -bonds and can be involved in noncovalent interactions similar to thoseformed by aromatic organic rings. Chelate rings can form XH. . .π and stacking interac-tions. XH. . .π interactions can be formed with organic molecules or ligands coordinatedto a metal. Stacking interactions can be formed with organic aromatic rings and with otherchelate rings.

The present chapter will discuss experimental observations and theoretical modeling ofthese types of noncovalent interactions.

10.2 Interactions with Organic π -Systems

10.2.1 Cation. . .π Interaction with Transition-Metal Complexes

Cation. . .π interactions are important noncovalent binding forces for artificial receptorsand in biomolecules. In this review, a special type of cation. . .π interaction where thecation is a cationic transition-metal complex will be described. In the case of an aromatic π -system, this interaction is termed a metal–ligand aromatic cation. . .π (MLACπ ) interaction.Figure 10.1 shows the [Co(NH3)6]3+ cation interacting with benzene. In the following,direct interactions of a π -system with a transition-metal atom, as exist in ferrocene andother sandwich compounds, where the π -system interacts with the metal’s d orbitals, arenot considered, as these interactions are better described as covalent interactions. Severalpapers dealing with the MLACπ interactions have been published in the last decade. Thesedescribe both experimental observations based on crystal-structure determinations as wellas quantum-mechanical calculations. This work will now be summarised.

Figure 10.1 Optimised geometries of the most stable conformers of [Co(NH3)6]3+-C6H6,[Co(NH3)6]3+-C2H4 and [Co(NH3)6]3+-C2H2 complexes.

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π -Systems of Transition Metal Complexes 245

Table 10.1 Bonding energies and geometric data for the most stable conformers of the[Co(NH3)6]3+-C6H6, [Co(NH3)6]3+-C2H4 and [Co(NH3)6]3+-C2H2 complexes

[Co(NH3)6]3+-C6H6 [Co(NH3)6]3+-C2H4 [Co(NH3)6]3+-C2H2

�E (kcal/mol)a −32.18 −17.02 −18.30The charge transfer 0.332 0.194 0.209Co-N (A) 2.02 2.02 2.03N-HINT (A) 1.04 1.04 1.04

2.68HINT

. . .� (A) 2.68 2.45 2.032.73

aBonding energies with BSSE correction; � is the centre of the benzene ring, of the double bond and the triple bond,respectively.

10.2.1.1 Predictions by Calculation

Cation. . .π interactions between a transition-metal complex cation and benzene, ethy-lene, and acetylene were predicted by theoretical calculations [1]. The calculations on the[Co(NH3)6]3+-C6H6, [Co(NH3)6]3+-C2H4 and [Co(NH3)6]3+-C2H2 systems demonstratethat a transition-metal complex cation can interact strongly with π -systems [1], results latersubstantiated by experiment [2]. The geometries of isolated molecules and of all conform-ers of the cation–π complexes were fully optimised and bonding energies were calculatedusing the B3LYP level of theory [3] with the 6-31G∗ basis set for the carbon, nitrogenand hydrogen atoms and LANL2DZ for the cobalt ion. The basis set superposition error(BSSE) correction was calculated [4]. Bonding energies and geometric data of the most sta-ble conformers of the complexes are collected in Table 10.1. The theoretical data predict thebonding energies for the most stable conformers of benzene, ethylene and acetylene com-plexes to be −32.18, −17.02 and −18.30 kcal/mol, respectively. In the most stable structureof the benzene π -complexes, three NH3 groups interact with the π -system, whereas in themost stable structure of the ethylene and acetylene π -complex, two NH3 groups interactwith the π -system (Figure 10.1). The charge transfer in the [Co(NH3)6]3+-C6H6 complexis considerably larger than in the case of [Co(NH3)6]3+-C2H4 and [Co(NH3)6]3+-C2H2

(Table 10.1).

10.2.1.2 Metal–Ligand Aromatic Cation. . .π Interactions in Crystal Structures

Cation. . .π (MLACπ ) interactions were identified [2] in crystal structures deposited in theCambridge Structural Database (CSD). Strong intermolecular interactions between ligandsin cationic metal complexes and aromatic rings of tetraphenylborate anion were identifiedusing the geometric criteria: distance between the metal and the centre of any aromaticring (�) of the tetraphenylborate anion less than dM = 5.5 A, the angle βM of the normalto the aromatic ring plane with the vector between the centre of the aromatic ring and themetal was less than 52◦, and the X–H. . .� angle was larger than 100◦ (Figure 10.2). Insome of the structures very short H. . .� distances were observed, below 2.30 A.

Based on the number of phenyl groups from the [BPh4]− anion interacting with thesame ligand, there are two types of interactions. In the first type, (I), one phenyl groupinteracts with one hydrogen atom of a ligand (Figure 10.3a), while in the second type, (II),

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dM

dX

Ω Ω

βX βα

βM

M X

H

M X

H

Figure 10.2 Geometric parameters describing MLXH. . .π interactions.

two phenyl groups interact with two hydrogen atoms of the same ligand (Figure 10.3b).For structures of type II, the distance between the hydrogen atoms and the centre of theappropriate phenyl ring are different. For one hydrogen atom, the distances are almost inall cases smaller than 2.3 A and smaller than the corresponding distance in the structures oftype I, which are in the range of 2.3 to 2.6 A. The shorter distances in the structures of typeII are the consequence of the simultaneous interaction of two phenyl rings of the [BPh4]−

anion with the transition-metal complex cation.

10.2.2 MLXH. . .π Interactions

The second type of interactions with organic π -systems are the MLXH. . .π interactions.The XH. . .π [5] and CH. . .π interactions [6] of coordinated ligands have been described incoordination complexes and organometallic compounds. In crystal structures, MLXH. . .π

interactions can be intramolecular [7] and also intermolecular as identified for acetylaceto-nate (acac) complexes [8]. Similar interactions have been also identified in metalloproteinstructures [9].

(a) (b)

Figure 10.3 Selected examples for MLACπ interactions with the [BPh4]− anion: (a) exampleof a type-I structure, as found in the crystal structure of NOKBAQ, where one phenyl groupinteracts with one hydrogen atom, and (b) example of a type-II structure, as found in thecrystal structure of ZUTNEH, where two phenyl groups interact with two hydrogen atoms ofthe same ligand.

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(a) (b)

Figure 10.4 Intramolecular MLXH. . .π interactions in complexes with (a) two bonds betweenthe metal atom and the aromatic group (structure HIPDUF, X = O), and (b) several bondsbetween the metal atom and the aromatic group (structure XOTQAJ, X = N).

10.2.2.1 Intramolecular MLXH. . .π Interactions

10.2.2.1.1 Intramolecular MLXH. . .π Interactions in Crystal Structures of Transition-Metal Complexes Crystal structures of metal complexes containing a phenyl ring anda transition metal were analysed for intramolecular XH. . .π interactions of a hydrogenatom bound to coordinated nitrogen or oxygen atoms. Interactions were observed in a largenumber of structures containing a variety of metal ions and ligands. Two examples areshown in Figure 10.4.

In some of the complexes, the aromatic group can be bound to the atom that is coordinatedto the metal ion (Figure 10.4a), that is, with two bonds between the metal atom and thearomatic group. In other complexes, the aromatic group can be separated from the metalcentre by several bonds (Figure 10.4b). In most of the complexes, the aromatic ring interactswith the hydrogen atom of another ligand. However, there are examples of aromatic groupsinteracting with a hydrogen atom of the same ligand.

Intramolecular metal-ligand OH. . .π (MLOH. . .π ) and metal–ligand NH. . .π

(MLNH. . .π ) interactions in transition-metal complexes between aqua or ammine ligandare described in more detail.

Employing the geometric criteria [7b]: (i) the distance between a hydrogen atom of aquaor ammine ligand and the centre of a C6-aromatic ring (Ω) was less than 3.5 A, and (ii)angles α ≥ 110◦ and β < 30◦ (Figure 10.2), intramolecular interactions were found in 38structures containing an aqua ligand and in 10 structures with an ammine ligand.

Most complexes that form intramolecular MLOH. . .π and MLNH. . .π interactions areneutral with only a few positively charged. The dominance of neutral complexes indicatesthat in neutral complexes, partial positive charge on hydrogen is sufficient to create an

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(a) (b) (c) (d)

Figure 10.5 Model systems for calculation of MLXH. . .π intramolecular interactions. Thecomplexes (a) and (b) were used as model systems for MLOH. . .π interactions, while (c) and(d) were used for MLNH. . .π interactions.

interaction with the π system of the aromatic ring. In positively charged complexes, largerpartial positive charge on the hydrogen atom increases its tendency to interact, while positivecharge does not decrease significantly the tendency of the aromatic ring to interact.

In compounds where the MLOH. . .π interactions are formed with two bonds betweenthe aromatic ring and the metal ion, geometric data indicate that interactions are not verystrong, probably due to steric constraints. Hence, the number of structures with more thantwo bonds between the aromatic ring and the metal centre is larger. In these structures,MLOH. . .π interactions are stronger, in accordance with fewer steric constraints. It isinteresting that in all analysed structures, the second hydrogen atom from a water moleculeis involved in hydrogen bonding with surrounding residues.

By contrast to the situation of MLOH. . .π interactions, when MLNH. . .π interactions areformed, a larger number of interactions with two bonds between metal and aromatic carbonare found. Thus, compared to the aqua ligand, the ammine ligands create more interactionswith shorter distances in structures with two bonds between metal and aromatic ring. Basedon this fact, the conclusion that could be drawn is that MLNH. . .π interactions in theseammine structures are more favourable.

10.2.2.1.2 Calculated Energies of Intramolecular MLXH. . .π Interactions DFT calcu-lations were performed on the model systems comprising Co(III) complexes as shown inFigure 10.5. The energies of the intramolecular MLXH. . .π interactions were evaluated byDFT calculations at the B3LYP [3] level of theory with the LANL2DZ basis set for thecobalt atom and the 6-31G(d,p) basis set for all other atoms.

For the aforementioned model systems, the energies of the conformations with andwithout MLXH. . .π interactions were calculated. The difference in the obtained energiescan be mainly attributed to the MLXH. . .π interactions. Based on these results, the energiesof the intramolecular MLOH/π interactions between the coordinated water molecule andthe aromatic ring were estimated to lie in the range −4.4 and −5.1 kcal/mol, whereas theenergies of the MLNH. . .π interactions were somewhat higher, that is, the range was −3.5/to −4.0 kcal/mol (Table 10.2).

10.2.2.2 Intermolecular MLCH. . .π Interactions Involving Acetylacetonate Ligands

10.2.2.2.1 MLCH. . .π Interactions of Acetylacetonate Ligands in Crystal StructuresThe crystal structures of metal complexes with coordinated acetylacetonate ligands were

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Table 10.2 Estimated energies of MLXH. . .π interactions in model systems

Structurea Type of interaction Nb �E (kcal/mol)

[CoCl2(NH3)2(C6H5NH2)(H2O)]+ MLOH. . .π 2 −5.1[CoCl(NH3)2(H2O)(Phe)]+ MLOH. . .π 4 −4.4[CoCl2(NH3)3(C6H5NH2)]+ MLNH. . .π 2 −3.5[CoCl(NH3)3Phe]+ MLNH. . .π 4 −4.0

aStructures are shown in Figure 10.5.bN is number of bonds between the aromatic ring and the metal centre.

screened to find metal–ligand CH. . .π (MLCH. . .π ) interactions involving acac CH-Hatoms [8], using the following geometric criteria: the distance between the H atom and thecentre of the phenyl ring to be shorter than 3.5 A, α > 110◦, β < 30 ◦ (Figure 10.6).

A large number of structures where an acac ligand, employing either CH-H ormethyl-Hatoms, is involved in interactions with a phenyl ring were found: 32 interactions with theCH-H atom and 584 interactions with a methyl-H atom. The significantly larger numberof interactions with methyl-H arises as a result of the larger number of available hydrogenatoms (there are two CH3 residues) and for steric reasons (the CH-H atom is often shieldedby the flanking methyl groups). Illustrations of examples of structures with CH-H andmethyl-H interacting are given in Figure 10.7.

An example of a MLCH. . .π interaction involving a CH-H atom with a phenyl ring isillustrated in Figure 10.7a. The data in Figure 10.8 show that the H. . .� distance is nevershorter than 2.7 A. The reason for this is relate to steric constraints. A phenyl ring cannotapproach closer to the CH-H atom because of steric clashes with either or both of the CH3

groups.The strength of the CH. . .π interactions depends on the partial positive charge on the

hydrogen atom. The acac ligand is negatively charged and it obtains some positive chargeby coordinating to a metal cation. Hence, the partial positive charge on the hydrogen atomof the CH-H group depends on the delocalisation in the chelate ring that in turn depends onthe nature of the bonding between the ligand and the metal. To determine the influence of

Figure 10.6 Geometric parameters used for the determination of the CH. . .π interaction. Thering is aromatic or a chelate ring with a delocalised π -system. � is the centroid of the ring, d isH. . .� distance, α is the C–H. . .� angle, β is the angle between the H. . .� line and the normalto the chelate ring, and R is the normal distance of H to chelate ring plane. A is any atom.

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(b)

(a)

Figure 10.7 (a) An example of MLCH. . .π interactions involving the CH-H atom of the acety-lacetonato ligand (as in the crystal structure of WANYUF) and (b) an example of MLCH. . .π

interactions with the CH3-H atom of the acetylacetonato ligand (as in the crystal structure ofSIBLIZ).

0

1

2

3

4

2.7–2.8 2.8–2.9 2.9–3.0 3.0–3.1Distance (Å)

3.1–3.2 3.2–3.3

Soft Lewisacids

Borderline Lewisacids

Hard Lewisacids

Nu

mb

er o

f st

ruct

ure

s

Figure 10.8 Histogram showing the distribution of geometric parameters for MLCH. . .π inter-actions involving the CH-H atom of acetylacetonato ligand for complexes with soft, borderline,and hard Lewis acids.

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2.5–2.60

2

4

6

8

10

12

14

16

18

20

2.6–2.7 2.7–2.8 2.8–2.9 2.9–3.0 3.0–3.1

Distance (Å)

3.1–3.2 3.2–3.3 3.3–3.4 3.4–3.5

Soft Lewis acids Borderline Lewis acids Hard Lewis acids

Nu

mb

er o

f st

ruct

ure

s

Figure 10.9 Histogram showing the distributions of the H. . .� distances for intermolecularCH3-H. . .phenyl interactions for neutral complexes with soft-, borderline-, and hard-Lewisacids.

the type of metal on the H. . .� distances, the complexes were classified into three groups:those containing soft metal ions, complexes with borderline metal ions, and complexes withhard metal ions. The histogram in Figure 10.8 shows the H. . .� distances for the differenttypes of metals in neutral complexes, since charges can also influence the H. . .� distance.The data show that complexes of soft acids have a tendency to make shorter H. . .�

distances than complexes with hard acids, which have a tendency to make longer H. . .�

distances. All distances in complexes with hard acids are above 3.0 A.Soft metal ions form bonds with substantial covalent character with the acac ligand. This

is because a soft metal ion enables the delocalisation of the negative charge of the ligand,increases the partial positive charge in the ligand, and therefore enables the formation ofstronger MLCH. . .π interactions with shorter H. . .� distances. In complexes with hardacids there is less covalent character of the metal–ligand bonds and less positive charge inthe ligand so that the H. . .� distances tend to be longer. The quantum-chemical calculationssupport this conclusion, as shown in Section 10.2.2.2.2.

An example of MLCH. . .π interactions formed between methyl-H and a phenyl ring isgiven in Figure 10.7b. The data in Figure 10.9 show that some H. . .� distances are quiteshort (2.4 A) in comparison with the distances for the MLCH. . .π interactions formedby the CH-H atom. This was attributed to the steric accessibility of the CH3 group.However, the number of structures with short distances is not large (Figure 10.9). Theinfluence of the hardness of the metal ion on the H. . .� distances in neutral complexesshows tendencies similar to the data for the interactions with the CH-H atom (Figure 10.8),that is, complexes with hard acids tend to make weaker interactions with longer H. . .�

distances (Figure 10.9).

10.2.2.2.2 Calculations of the MLCH. . .π Interactions Involving AcetylacetonateLigands The model systems for calculations of the interactions of the CH-H and methyl-Hatom of acac complexes with aromatic rings are illustrated in Figure 10.10. Single-pointcalculations for the CH-H atom show that for distances below 2.6 A, the interaction is

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Figure 10.10 (a) The model system for the single-point MP2 calculations of the CH-H. . .π

interactions and (b) CH3-H. . .π interactions.

repulsive. This is in agreement with experimental data where no interaction with a distanceless than 2.7 A was observed (Figure 10.8). The minimum energy corresponds to a distanceof 3.1 A with an interaction energy of 1.78 kcal/mol.

The most favourable interaction involving methyl-H is 1.65 kcal/mol at a distance of2.65 A. This is in agreement with the short distances observed in the experimental data.The large number of interactions observed at the distances above 3.0 A (Figure 10.9) isa consequence of the interactions with the phenyl ring not only with one but with twoor three hydrogen atoms from the same CH3 group, as for example in Figure 10.7b. Thisalso explains the larger calculated interaction energy for the CH-H atom compared to themethyl-H atom.

10.2.2.2.3 Metal-Dependent Capacity for MLCH. . .π Interactions of AcetylacetonateLigands As discussed above, the H. . .� distances in the crystal structures indicate ametal-dependent capacity for the formation of CH. . .π interactions by the acac ligand. ForMLCH. . .π interactions, the H. . .� distances are shorter in complexes with soft metals (Fig-ures 10.8 and 10.9), as discussed above. To support these observations, quantum-chemicalcalculations on interacting energies were performed. It was not possible to compare hardmetal ions with the soft metal ions as they do not form the same type of complexes, namelysoft metal ions prefer square-planar complexes while hard metal ions tend to form octahedralcomplexes. The calculations were performed on neutral complexes: [M(acac)2], complexesof borderline, Ni(II), and soft acids, Pd(II) and Pt(II), and [M(acac)(en)] complexes of softacids, that is, Rh(I) and Ir(I).

Using the MP2 level of theory, the interacting energies between benzene and acac-boundCH-H and methyl-H atoms, corrected for BSSE, were calculated at the H. . .� distancesof 3.1 and 2.6 A, respectively, since it was found that the strongest interactions occur atthose distances (Section 10.2.2.2.2). The increase in the interacting energies from Ni to Ptin the [M(acac)2] series of complexes, and from Rh to Ir in the [M(acac)(en)] series forboth the CH-H and methyl-H atom (Table 10.3) are in agreement with the experimentaldata (Figures 10.8 and 10.9). The results also show that the oxidation state of the metaland other ligands in the complex exert a strong influence on the interacting energies. Thecalculated interaction energies are larger for the [M(acac)2] series than for [M(acac)(en)]complexes (Table 10.3).

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Table 10.3 Calculated energiesa of CH. . .π interactions (kcal/mol) for [M(acac)2] (M = Ni,Pd, Pt) and [M(acac)(en)] (M = Rh, Ir)

Type of interaction/M Ni Pd Pt Rh Ir

CH(acac). . .benzene 2.12 2.26 2.34 1.44 1.51CH3(acac). . .benzene 1.64 1.76 1.89 0.64 0.73

aCalculations were performed using the MP2/B3LYP level of theory with LANL2DZ and 6-31G∗∗ basis sets (see text).

10.2.2.3 MLXH. . .π Interactions in Metalloproteins

The structures of metalloproteins included in the Protein Data Bank were screened [9] forMLXH. . .π interactions occurring between metal-coordinated ligands and the aromaticresidues of phenylalanine, tyrosine, and tryptophan using the following geometric criteria:the distance between the metal and the centre of the aromatic ring, dM, was less than 5.5 A,and the angle, βM, between the normal to the aromatic ring plane with the distance vectorbetween the centre of the aromatic ring and the metal was less than 52◦ (Figure 10.2).

It was found that specific MLXH. . .π interactions between a coordinated ligand and thearomatic ring of phenylalanine, tyrosine and tryptophan are present in a number of crystalstructures of metalloproteins [9], involving different metals and a variety of ligands. Amongthe coordinating ligands are amino acids, such as asparagine, aspartate, glutamate, histidineand threonine, but also water and substrates like ethanol form this type of interaction.

As in the case of metal complexes, partial positive charge on the ligand, induced by thecharge on the metal, is situated on the hydrogen atom interacting with the aromatic ring. Inalmost all proteins investigated, the positions of hydrogen atoms were not available fromthe crystal structure. Hence, the geometry was characterised by the distance dX (betweenthe centre of the aromatic ring and the hydrogen carrier atom closest to the ring centre)and the angle βX (between the corresponding distance vector dX and the normal of thearomatic ring plane (Figure 10.2). For the interactions considered, the distance dX variedbetween 3.09 and 4.41 A. The angle βX was often less than 30◦ but, was sometimes as largeas 45◦.

In the structures of Mn and Fe superoxide dismutases (SOD), the distances, dX, betweenthe aromatic group of tryptophan involved in the MLXH. . .π interaction and the oxygenatom of coordinated water range from 3.12 to 3.57 A. This is generally shorter than thecorresponding distances between the two water oxygen atoms and Phe14 in the bacterialchemotaxis protein (CheY), that is, 3.50 and 4.12 A. Based on the comparison of the dis-tances involving the phenylalanine and tryptophan, it seems that tryptophan enters into thestrongest interactions with coordinated ligands. This is in agreement with calculations show-ing that tryptophan interacts more strongly than benzene in cation—π interactions [10].

The interaction will generally be stronger, the fewer the number of bonds between theatom coordinated to the metal and the interacting hydrogen atom. Water is the ligand withthe smallest number of bonds between the atom coordinated to a metal cation and theinteracting hydrogen atom; it involves just a single O–H bond, so that interactions of waterare anticipated to be the strongest. The data show that the shorter dX distance, of 3.12 A,occurs between water and tryptophan in a Fe(III) SOD. A representative structure witha MLXH. . .π interaction, that is, a structure of superoxide dismutase (SOD, PDB code1avm/A) [11] is shown in Figure 10.11.

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Figure 10.11 An example of the MLXH. . .π interaction (X = O) in superoxide dismutase(SOD, PDB code 1avm/A). All ligands coordinated to the metal centre and all aromaticresidues relevant for MLOH/π interaction are displayed. The putative O. . .π interactions areshown as dashed lines.

In the iron(III) SOD structure, Figure 10.11, determined to a resolution of 1.55 A,there are two tryptophans close to the metal centre, and each appears to be involved in aMLOH. . .π interaction. There are four amino acids coordinated to the iron: three histidines(His27, His75 and His165) and one aspartate (Asp161). The other two ligands are water andazide. The MLOH. . .π interaction with the shortest distance corresponds to the interactionof the coordinated water with the five-membered ring of Trp163, with dX = 3.35 A. Thesecond MLXH. . .π interaction appears to occur with the six-membered ring of Trp126that in turn interacts with the β-CH-H atom of Asp161 (Figure 10.11). While the distancebetween the centre of the six-membered ring to the β-CH-H atom is 4.77 A, it is noted thatthe O atom approaches one of the C atoms at a distance of only 3.65 A, suggestive of aninteraction with one of the delocalised C–C bonds.

Importantly, the two tryptophans involved in MLXH. . .π interactions in SOD are amongthe twelve residues of FeSOD and MnSOD that are strictly conserved. These tryptophanresidues probably play an important role in the enzymatic reaction mechanism of SOD,and the MLXH. . .π interactions may be important for their function.

In a similar way, the MLXH. . .π interactions can be important for the function of alcoholdehydrogenase Zn-metalloenzyme (ADH). The results of ab initio calculations showed thatthe stabilisation energy provided by the presence of phenylalanine (Phe93) to the activebinding site of ADH is about 3 kcal/mol [12].

10.3 Interactions with π -Systems of Chelate Rings

Planar chelate rings with delocalised π -bonds are π -systems specific to transition-metalcomplexes and organometallic compounds. In chelate rings, the metal atom is a member

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of the ring and can be involved in the π -delocalisation. There are indications that chelaterings can be aromatic and a term “metalloaromaticity” has been coined to categorise thisphenomenon [13]. The aromaticity of the chelate rings has been investigated theoreticallyby calculating nucleus-independent chemical shifts (NICS) [14]. The calculations show thatsome of the chelate rings can be aromatic, while many of them are not [15–18]. A specialexample of chelate rings, metallabenzenes, may be aromatic [16] while the aromaticityof some dithiolate chelate rings is marginal [17]. Calculations on metal complexes of o-benzoquinone diimine (bqdi, known to function as a very strong acceptor of π -electrondensity in certain complexes) chelate rings show that only the Ru2+-bqdi chelate ringhad large negative NICS values, indicating aromaticity [18]. A number of diimine chelaterings can be generally characterised as nonaromatic [17] as too can some acac chelaterings with various metals [18]. However, a number of planar chelate rings that are notaromatic can still form noncovalent interactions similar to organic aromatic molecules,as these planar chelate rings have regions of delocalised π -electron density. Structures inthe CSD were analysed and based on this analysis, we described CH. . .π and stackinginteractions of chelate rings [8,18–28]. Other groups have reported CH. . .π [29–32] andstacking interactions [34–37] of chelate rings in a number of crystal structures.

10.3.1 CH. . .π Interactions with π -Systems of Chelate Rings

10.3.1.1 CH. . .π Interactions with π -System of Chelate Rings in Crystal Structures

Specific CH. . .π interactions between CH-H atoms and the π -systems of chelate rings wereobserved by analysing the crystal structures of transition-metal complexes incorporated inthe CSD. A total of 5091 structures containing almost planar chelate rings with delocalisedπ -bonds were found [19]. Among these structures, there were 1265 with weak CH–π

interactions (α > 110◦, β < 16◦, and H···� distance < 3.0 A; α, β, and � are definedin Figure 10.6) and 151 structures considered to have very strong interactions (α > 145◦,β < 12◦, and H···� distance < 2.6 A). In a number of structures there were more thanone interaction per molecule. The geometric parameters characterising some examples areshown in Table 10.4, all of which feature a very short H···� distance (H···� < 2.6 A), withthe H atom almost centred over the chelate ring centre (β < 12.5◦) and the X-H···� anglenot too bent (α > 145◦). From a geometric point of view, interactions in these structures aresimilar to strong CH. . .π interactions in organic molecules. A representative example fromTable 10.4 is illustrated in Figure 10.12. Data in Table 10.4 show that these interactionsoccur in metal complexes, including organometallic compounds, of various metals andchelate rings. Of the interactions included in Table 10.4, some of the C donor atoms are sp3

hybridised, while others are sp2 hybridised. All interactions are intermolecular and occurbetween two equivalent neutral complexes.

10.3.1.2 CH. . .π Interactions of Aryl-C–H with Chelate Rings of Acetylacetonate

10.3.1.2.1 CH. . .π Interactions of aryl-C–H with the Acetylacetonate Chelate Ring inCrystal Structures As mentioned above, one of the chelate rings often featuring a delo-calised π -system is the one formed by acac and it can participate in CH. . .π interactions.The interactions that were specifically studied were those with the C6 group as the hy-drogen atom donor [8]; an example is shown in Figure 10.13. In all, 41 intermolecular

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Table 10.4 Selected crystal structures of metal complexes where a C-H group interacts withthe π -system of a chelate ringa

Refcode α (◦)b β (◦)c H. . .�d (A) C. . .� (A) Chelate ring De

BEPZOL 176 3.8 2.48 3.56 CuNCNCN Csp3

COTQAD 155 1.2 2.38 3.39 ZrNCCCN Csp2

COTQEH 155 0.6 2.36 3.38 HfNCCCN Csp2

FEBTAH 168 1.4 2.39 3.45 NiNCCCN Csp3

GORYER 177 1.6 2.46 3.54 MnOCCCN Csp3

GOXDAY 151 5.0 2.48 3.46 FeNCCCN Csp2

HALWAS 147 3.9 2.40 3.36 ZnNCCCN Csp2

HAMDAA 149 3.7 2.40 3.38 ZnNCCCN Csp2

HAMGUX 149 4.7 2.46 3.43 ZnNCCCN Csp2

JOXGIM 158 6.2 2.59 3.68 CuOCCCN Csp3

LICSEV 154 4.7 2.31 3.32 CuNCCCN Csp3

LOFZOV 151 3.4 2.36 3.35 NiOCCCN Csp2

MAMTOJ 166 5.3 2.46 3.52 CoOCCCN Csp2

MAMTUP 16 1.0 2.32 3.38 CuOCCCN Csp3

PIQZUK 175 4.7 2.45 3.53 ReOCCCN Csp2

SUMXIH 146 5.1 2.50 3.45 FeNCCCN Csp3

VAFQIC 174 3.6 2.30 3.38 ZnNCCCN Csp2

YINLUC 157 3.1 2.54 3.60 NiOCCCO Csp2

ZAXDOR 156 5.3 2.30 3.31 ZnOCCCN Csp3

ZONQOI 149 4.4 2.34 3.32 CoNCCCN Csp3

ZOPYUY 157 12.4 2.34 3.37 CuOCCCN Csp3

aData for normalised H-atom positions.bα is the C-H···� angle (Figure 10.6).cβ is the angle between the normal of the chelate ring and the �···H line (Figure 10.6).dH···� is the distance between the normalised H-atom position and the centre of the ring (Figure 10.6).eCsp2 is bonded to two nonhydrogen atoms. Csp3 is bonded to two or three nonhydrogen atoms.

Figure 10.12 Perspective view of BEPZOL as a selected example for a CH. . .π interactionwith the π -system of the chelate rings.

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Figure 10.13 Crystal structure of XIBRUV representing an example of CH. . .π interactionswith the acac chelate ring acting as the hydrogen-atom acceptor.

interactions in 36 crystal structures included in the CSD were identified satisfying the ge-ometric criteria: d < 3.5 A, α > 110◦, β < 30◦ (Figure 10.6). The shortest H. . .� distancesare close to 2.4 A, similar to those of the MLCH. . .π interactions formed by acac-methyl-H(Section 10.2.2.2.1).

The complexes with CH. . .π interactions were classified in three groups: complexes withmetals that are soft, borderline and hard acids. In Figure 10.14, histograms for square-planar(soft and borderline metal ions) and octahedral complexes (borderline and hard metal ions)are presented. As anticipated, the H. . .� distances are shorter in the square-planar structuresthan in the octahedral complexes. The H. . .� distances for square-planar complexes withborderline metal ions are below 2.7 A (with one exception), while they are over 3.0 Ain octahedral complexes. Data in the histograms for square-planar complexes indicate thatthere is some tendency of complexes with borderline metal ions to form stronger interactionsthan complexes with soft metal ions (Figure 10.14). The complex with an exceptionallylong distance has a lot of steric constraints. Hence, it seems that better delocalisation of thenegative charge in complexes with soft metal ions does not favour the interaction.

In octahedral complexes, a chelate ring is less accessible for the CH. . .π interaction andsteric constraints have greater influence on the formation of such interactions. Hence, onecannot expect to observe the influence of the type of metal on the geometric parameters, ascan be seen in Figure 10.14.

Most of the complexes forming CH. . .π interactions are neutral and the interactionusually occurs between two equivalent complexes. There are only two examples of inter-actions of positively charged and three for negatively charged complexes. Significantly, theinteraction also includes complexes with negative charge, while in the case of MLCH. . .π

they were absent (Section 10.2.2.2.1). In all three examples the C6 ring is part of a cation.In the two structures with positively charged complexes, the interactions occur between

two equivalent molecules. In these cases the H. . .� distances are quite short, 2.66 and 2.80A. It seems that the positive charge in the complex does not significantly affect the capabilityof the π -system of the chelate ring to participate in the CH. . .π interactions because thepositive charge is located mainly on the hydrogen atoms of the acac chelate ring. On the

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2.4–2.5

0

1

2

3

4

5N

2.5–2.6 2.6–2.7 2.7–2.8 2.8–2.9 2.9–3.0

Distance (Å)

Square-planar acetylacetonato complexes

3.0–3.1 3.1–3.2 3.2–3.3 3.3–3.4 3.4–3.5

2.4–2.5 2.5–2.6 2.6–2.7 2.7–2.8 2.8–2.9 2.9–3.0 3.0–3.1 3.1–3.2 3.2–3.3 3.3–3.4 3.4–3.50

1

2

3

4N

Distance (Å)

Octahedral acetylacetonato complexes

Soft Lewis acids Borderline Lewis acids Hard Lewis acids

Figure 10.14 Histograms showing the distributions of the H. . .� distances for the intermolec-ular C–H. . .chelate interactions in neutral square-planar and octahedral complexes.

other hand, a larger partial positive charge on the hydrogen atoms of C6 ring favours theCH. . .π interaction, so the interaction is quite strong.

10.3.1.2.2 Calculated Energies for CH. . .π Interactions of Benzene with Chelate Ringsof Acetylacetonate The calculation of the interaction of the hydrogen atom of the C6

ring with the π -system of the chelate ring was performed on the model system shown inFigure 10.15 [8]. The calculations were performed at the B3LYP level of theory, with theLANL2DZ basis set for the metal atom and the 6-31G

∗∗basis sets for other atoms, using

the Gaussian98 programe.The most favourable interaction occurs at a H. . .� distance of 2.6 A with energy

of 1.62 kcal/mol. The minimum distance is almost the same as the minimum for the

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Figure 10.15 The model system for the calculations of the CH. . .π interactions.

MLCH. . .π interactions formed by methyl-H (Section 10.2.2.2.2) and the interacting en-ergies at the minima are comparable, indicating that the nature of these interactions aresimilar.

10.3.1.2.3 Metal-Dependent Capacity for the Formation of the CH. . .π Interaction withthe Chelate Rings of Acetylacetonate As discussed in Section 10.3.1.2.1, data from crystalstructures having CH. . .π interactions involving the π -system of the chelate ring in thesquare-planar complexes indicate that the H. . .� distances are shorter with borderline metalions than in the complexes with soft metal ions (Figure 10.14). The interaction energies ofthe benzene-C–H with the π -system of the chelate ring were calculated for the [M(acac)2](M = Ni, Pd and Pt) and [M(acac)(en)] (M = Rh and Ir) series of complexes at a H. . .�

distance of 2.6 A because the strongest interaction is at that distance. The calculatedenergies are in the range 1.58 to 2.63 kcal/mol (Table 10.5). The interacting energies ofthe [M(acac)2] complexes for a borderline acid, Ni(II), and soft acid ions, Pd(II) and Pt(II),are very similar, while the energies for the [M(acac)(en)] complexes of the two soft acids,Rh(I) and Ir(I), are quite different [8].

These data indicate that the interacting energies of benzene with the π -system of the acacchelate ring do not depend on the hardness of the metal, as was the case for interactionswhere acac was the hydrogen-atom donor (Section 10.2.2.2.1). A comparison of the resultsfor the [M(acac)2] complexes with those for the [M(acac)(en)] complexes show that theCH. . .π interactions with acac chelate ring are stronger for the latter (Table 10.5).

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Table 10.5 Calculated energiesa of CH. . .π interactions (kcal/mol) for [M(acac)2] (M = Ni,Pd, Pt) and [M(acac)(en)] (M = Rh, Ir)

M Ni Pd Pt Rh Ir

CH. . .π (benzene-chelate) 1.60 1.58 1.59 2.42 2.63

aCalculations were done using the MP2//B3LYP level of theory with LANL2DZ and 6-31G∗∗ basis sets (see text).

10.3.1.2.4 The Nature of CH. . .π Interactions with the Chelate Rings of Acetylaceto-nate In order to get more insight about the nature of the CH. . .π interactions involvingchelate rings and their similarity with interactions involving organic aromatic rings, calcu-lations were performed investigating the interaction of a benzene–CH– with the π -systemof an acac chelate ring in a Ni(II) complex, Figure 10.15, and on the CH. . .π interactionformed between two benzene molecules [20]. Single-point calculations for the CH. . .π

interaction were performed for distances ranging from 2.1 to 3.5 A between the centreof the ring (�) and the hydrogen atom of the benzene molecule. For the calculations,B3LYP and MP2 levels of theory were used. The MP2 data were corrected for BSSE.A comparison of the complementary B3LYP and MP2 levels of theory enabled a con-clusion to be drawn about the nature of the CH. . .π interaction with the chelate ring.Namely, the B3LYP level of theory can provide information about the electrostatic com-ponent but not the dispersion component, while the MP2 level of theory can identify thedispersion component.

The interacting energies with B3LYP level of theory for the benzene. . .chelate andbenzene. . .benzene CH. . .π interactions are quite similar. At the minima, the energies are−0.68 and –0.64 kcal/mol, respectively (Table 10.6). The negative energies obtained forB3LYP level of theory show that there is an attractive electrostatic interaction for both Ni(II)-acac chelate and benzene ring systems. The results for benzene are in agreement with theresults that show that the T-shaped benzene dimer is stabilised by an attractive electrostaticinteraction [33]. The somewhat larger B3LYP energy for the Ni(II)-acac chelate ring showsthat the electrostatic components are stronger. The interacting energies with the MP2level of theory for the benzene. . .chelate and benzene. . .benzene CH. . .π interactions areagain very similar; at the minima the energies are −1.62 and −1.58 kcal/mol, respectively(Table 10.6).

A comparison of the energies obtained with the MP2 and B3LYP levels of theory indicatesthat the dispersion component is a major component contributing to the interacting energy

Table 10.6 Interaction energies at the minima for benzene–chelate and benzene–benzeneinteractions using MP2 and B3LYP levels of theory

Benzene–chelate Benzene–benzene

Theory �E (kcal/mol) Distance (A) �E (kcal/mol) Distance (A)

B3LYP −0.68 2.8 −0.64 2.9MP2 −1.62 2.6 −1.58 2.7

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and that the CH. . .π interaction involving the Ni(II)-acac chelate ring has almost the samecontribution of dispersion component as the CH. . .π interaction of two benzene molecules.Again, the results for benzene are in agreement with previous results that show that for T-shaped benzene the major source of attraction is dispersion interaction [33]. For both levelsof theory, the calculated interaction energy is larger for benzene–chelate interaction andthe distances are shorter. Both energies and distances indicate that the CH–π interactioninvolving the Ni(II)-acac chelate ring in the model system is somewhat stronger thanthe CH. . .π interaction involving two benzene rings. Very accurate calculations on theT-shaped benzene dimer show that the interacting energy is −2.39 kcal/mol. Based onthat, the CH. . .π interaction involving the Ni(II)-acac chelate ring in the model system isestimated to be about −2.5 kcal/mol.

It is well known that CH. . .π interactions of organic aromatic molecules are very impor-tant in different molecular systems, including biochemical systems. Based on the similarinteracting energies of π -systems of acac chelate and benzene rings, CH. . .π interactionsmay very well be as important as their all-organic counterparts.

10.3.1.3 CH. . .π Interactions in Metal-Porphyrin Complexes

The porphyrin molecule and derivatives possess four pyrrole rings with delocalised π -systems. When a porphyrin is coordinated as a tetradentate ligand to a metal centre,additional rings with delocalised π -bonds are formed, that is, four six-membered chelaterings. These chelate ring π -systems can potentially be involved in CH–π interactions.Such interactions have been identified in the crystal structures of transition-metal com-plexes [21,22] and in structures of porphyrin-containing proteins [23]. A summary ispresented below.

10.3.1.3.1 CH. . .π Interactions in Crystal Structures of Metal–Porphyrin Complexes Incrystal structures of transition metal–porphyrin complexes included in the CSD, 655 specificCH. . .π interactions, formed between a C–H atom and the π -system of the porphyrin chelaterings, were found using the geometric criteria: the distance between an hydrogen atom andthe centre of the six-membered ring (�) was shorter than 3.0 A, α > 110◦ and β < 16◦

(Figure 10.6) [22]. In a number of cases, more than one interaction is observed in the samecrystal structure. Among the 655 “hits”, 411 are intermolecular and 244 intramolecular.Intermolecular CH. . .π interactions are possible only in metal complexes with at least onevacant axial position. Histograms showing the distribution of the geometric parameters forinter- and intramolecular CH. . .π interactions are presented in Figure 10.16, and significantdifferences are evident.

The peaks of the distribution of distances are in the regions of 2.2–2.4 and 2.6–2.8 A,for intra- and intermolecular interactions, respectively. The β angles for the intramolecularinteractions occur in the full range 0–16◦, while for the intermolecular interactions, thefull region is again covered but with a tendency for wider angles. For a large number ofintramolecular interactions, the α angle is small, that is, between 110 and 120◦, while forthe intermolecular interactions, the peak of the distribution is at 140–150◦. The differencesin the distributions of geometric parameters for intra- and intermolecular interactions area consequence of geometric constraints in the intramolecular interactions. In most of the

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Figure 10.16 Histograms showing the distribution of geometric parameters (dx, β and α) forintermolecular (a) and intramolecular (b) CH. . .π interactions. For graphs of the intramolecularinteractions, bars in black are data for all intramolecular interactions and those in gray are datafor intramolecular interactions after removing structures containing axially coordinated pyridineand similar types of ligands.

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Figure 10.17 The model system for the calculations of the C–H. . .chelate interactions incomplexes of porphyrin.

intramolecular interactions the hydrogen atom that interacts with the chelate ring is part ofthe axial ligand.

10.3.1.3.2 Calculated/Energies for CH. . .π Interactions in Metal–Porphyrin ComplexesThe interaction energy for CH. . .π interactions with a chelate ring of a coordinated por-phyrin was determined by DFT calculations on the [Cu(porphyrin)]–C2H2 model systemwhere the hydrogen atom of the acetylene interacts with a chelate ring (Figure 10.17) [22].Single-point calculations have been performed for different H···� distances and the mini-mum at 2.6 A corresponds to an interaction energy of 1.58 kcal/mol. The curve shows thatthe attractive interactions persist even at distances longer than 3.0 A.

For systems with dispersion forces, DFT calculations give reasonable geometries.Hence, one can assume that calculated DFT geometries are very accurate, while calculatedenergies are somewhat underestimated. Based on the energies obtained for the C2H2–C6H6

system, the estimated interaction energy for the [Cu(porphyrin)]–C2H2 system is about3.0 kcal/mol [22].

The DFT calculated distance for the most stable structure is in good agreement with thedistances of the intermolecular interactions found in the crystal structures (Figure 10.16).The calculations show that the energies of interactions with distances shorter than 2.6 A aresmaller. Hence, shorter distances in the intramolecular interactions are not due to strongerinteractions, but to geometrical constraints, as mentioned above. An illustrative example isa structure with an axially coordinated pyridine (Figure 10.18).

An example featuring an intermolecular CH. . .π interaction is found in the structure ofthe copper(II) complex (KIKGIU). The interaction exists between a six-membered ring anda C–H atom of a chloroform molecule (Figure 10.19). The hydrogen atom is almost ideallyorientated towards the centre of the six-membered ring at a very short H. . .� distance of 2.23A, with β = 11.6◦ and C–H. . .� = 177◦. This very favourable interaction is a consequenceof the freedom of the chloroform molecule, as an independent solvent molecule, to orientitself in an optimal fashion.

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Figure 10.18 Interactions of α-hydrogen atoms of pyridine with chelate rings of a porphyrinatoligand, as in the structure of CPOEFE10.

10.3.1.3.3 CH. . .π Interactions in Porphyrin-Containing Proteins A search of the Pro-tein Data Bank for porphyrin-containing proteins revealed that the π -system of everyporphyrin ring is involved in CH. . .π interactions, with the majority of the porphyrins hav-ing several interactions [23]. Both five-membered pyrrole rings and six-membered chelaterings are involved in CH. . .π interactions. It is noted that the number of interactions involv-ing five-membered rings is greater than the number of interactions with six-membered rings.

10.3.2 Stacking Interactions Involving π -Systems of Chelate Rings

An analysis of the crystal structures archived in the CSD showed that planar chelate ringscan form stacking interactions with C6-organic aromatic rings (chelate-phenyl) and withother chelate rings (chelate–chelate). These types of interactions are now summarised.

10.3.2.1 Stacking Interactions between Chelate and C6 Aromatic Rings

Stacking interactions of chelate with C6 rings were first recognised in the crystal structuresof square-planar Cu(II) complexes based on a survey of the CSD [24]. Subsequently, it was

Figure 10.19 Two approximately orthogonal projections highlighting Cl3C–H interactionswith a porphyrin derivative as in the structure of KIKGIU.

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P1

P2

Ct βD

Figure 10.20 Geometric parameters describing the interaction of square-planar complex witha C6 ring. The chelate ring has a delocalised π -system and can be five- or six-membered. D isdistance between the metal and the closest aryl carbon, Ct is the distance between the centresof the chelate and C6 rings, P1 is mean plane of metal and coordinated atoms, P2 is meanplane of the C6 ring, and β is the angle between the normal to the C6 ring and the vector thatconnects the centres of the chelate and C6 rings.

shown that square-planar complexes of various metals can form these interactions and thatthese interactions are in fact ubiquitous [25,26].

Interactions between chelate and C6 rings in crystal structures of square-planar com-plexes of various transition metals were recognised by comparing geometric parameters(Figure 10.20) for complexes with and without chelate rings. Structures were analysedwhere only complex molecules were present in the asymmetric unit. In all, 604 crystalstructures were identified and of these 285 contained a planar chelate ring and 319 did nothave a planar chelate ring. There was a broad representation of metal atoms amongst thestructures. All ligands that formed the chelate rings have π -bonds prior to coordination tothe metal atom and could potentially be involved in stacking interactions with C6 rings.

An analysis of crystal structures of complexes with and without planar chelate ringsshowed that the presence of a planar chelate ring has a clear influence upon the metal–C6

carbon distance. Histograms that show the distributions of the number of structures versusthe parameter �, which is the difference between the sum of the van der Waals radii and thedistance between the metal and the closest C6 carbon (D) (Figure 10.20) (� = �vdW −D), are presented in Figure 10.21. The parameter � enabled the use of a common criterionfor all metal atom types. In the complexes with planar chelate rings, a larger percentage ofstructures have a metal–carbon distance below the sum of the van der Waals radii (� > 0)than in the complexes without planar chelate rings.

Figure 10.21 Histograms showing the distribution of parameter � (A) in crystal structureswith (a) and without (b) planar chelate rings.

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Figure 10.22 Histograms showing the distribution of the dihedral angle θ (◦) between themean plane of the C6 ring (P2) and the mean plane of the metal with coordinated atoms (P1)with (a) and without (b) planar chelate rings.

The other important geometric parameter is the dihedral angle θ , that is, the anglebetween the mean plane of the C6 ring (P2) and the mean coordination plane comprisingthe metal and atoms forming the chelate (P1) (Figure 10.20). The distribution of the θ anglevalues in structures with and without a planar chelate ring is shown in Figure 10.22. Inthe structures with a planar chelate ring, the C6 ring has a large tendency to be orientedapproximately parallel to the mean coordination plane (Figure 10.22a). By contrast, inthe structures without a planar chelate ring, the C6 ring has a tendency to be orientedorthogonal to the coordination plane (Figure 10.22b). The short metal–carbon distancesand approximately mutual parallel orientation of the C6 and planar chelate rings indicatearyl. . .chelate stacking interactions.

One of the consequences of the overlapping of two rings is correlation of the distancesbetween the centres of the chelate and C6 rings (Ct) and metal–carbon distances (D)(Figure 10.20), as shown in Figure 10.23. The data show that when the distance betweenthe centres of these two rings is short then the metal–carbon distance is also short. Thiscorrelation can only exist if the chelate and C6 rings tend to overlap.

Figure 10.23 (a) Plot of the distances between the centres of the chelate and C6 rings Ct (A)versus the metal–carbon distance D (A). (b) Plot of the dihedral angle ϕ (angle between themean planes of chelate and C6 rings) versus angle β (angle between the normal to the C6 ringand the line connecting the centres of the two rings (Figure 10.20)).

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Figure 10.24 Plot of the dihedral angle ϕ (the angle between the mean planes of chelate (P1)and C6-aryl (P2) rings) versus the distance between the centres of the two rings (Figure 10.20).

The positions of C6 rings with respect to chelate rings show that in all cases with smalldihedral angles ϕ (the angle between the mean planes of chelate and C6 rings) the two ringsare in mutual slipped-parallel orientations (offset face to face). The values of the dihedralangle ϕ plotted versus angle β (Figure 10.20) are shown in Figure 10.24. The plot showsthat when the two rings are parallel (small dihedral angle ϕ), the β angle has, in most cases,values between 15 and 35◦. This indicates a slipped-parallel orientation of the two rings.The position of the peak distribution of the metal–aryl carbon distances in the region belowthe sum of van der Waals radii, the positive correlation of metal–carbon distances with thedistances between centres of chelate and C6 rings, along with the mutual slipped-parallelorientation of two rings, show that stacking interactions between C6 and chelate rings canexist in square-planar complexes of various transition metals. The data also show that thegeometry of the stacking interaction between C6 and chelate rings is similar to the geometryof the stacking of two benzene rings [33], indicating that chelate rings can behave similarlyto organic aromatic rings.

10.3.2.2 Comparison of Stacking and CH. . .π. . .Interactions betweenChelate and Aryl Rings

Since chelate rings can form CH. . .π (Section 10.3.1) and stacking interactions (Sec-tion 10.3.2.1) with aryl rings, it was thought of interest to analyse the tendency for stackingversus CH. . .π interactions [27]. The crystal structures of square-planar transition-metalcomplexes deposited in the CSD with close contacts between planar chelate rings andC6-aryl rings were analysed. There were 268 structures where the distance between thecentres of the two rings was shorter than 5.5 A. It is interesting to note that in most of thesesquare-planar complexes the chelate rings are fused with aromatic or other π -delocalisedchelate rings. Noncovalent interactions in these structures were analysed.

The diagram plotting ϕ, the angle between the mean planes of the chelate (P1) andC6-aryl (P2) rings, versus Ct (Figure 10.20) is presented in Figure 10.24. As mentionedin Section 10.3.2.1, and as is evident from Figure 10.24, in most of the structures theangle ϕ is small and the two rings are parallel as a consequence of the π. . .π stacking

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interactions. However, the data in Figure 10.24 show that there are numerous structureswith quite large ϕ angle, that is, up to 90◦, that correspond to an orthogonal orientation ofthe two rings. The orthogonal orientation indicates the possibility for CH. . .π interactionsand, indeed, CH. . .π interactions were indentified in a significant number of structures. Ina small number of structures CH. . .X (X = O, N, S) interactions were found. There werealso a small number of structures where none of the above interactions were recognised.

It was considered that two rings form stacking interaction when ϕ > 20◦, the dihedralangle between the mean planes of chelate (P1) and C6-aryl (P2) rings (Figure 10.20). Thereis a possibility that stacking occurs when a C6-aryl ring overlaps with a chelate ring or withthe ring that is fused to the chelate ring: CH. . .π interactions occur with the chelate ring ifthe distance between the centres of the chelate and C6-aryl rings is shorter than the distancebetween the centres of the fused and C6-aryl ring, while CH. . .π interactions occur with thefused ring if the distance between the centres of the fused and C6-aryl rings is shorter. Thedata related to the intermolecular CH. . .π interaction were analysed using the structuralparameters defined in Figure 10.6. An intermolecular interaction was assigned when α >

110◦, β < 30◦ and the normal distance of the hydrogen to the chelate ring plane was shorterthan 3 A, that is, R < 3 A. To investigate the geometry of the CH/X contact present in theretrieved data set, a search was carried out for acceptor species X = N, O and S that wasthought significant if the H. . .X distance was less than 3.0 A.

The number of structures with different types of noncovalent interactions is presentedin the histogram shown in Figure 10.25. The data show that planar chelate rings can beinvolved in stacking and CH. . .π interactions with organic aryl rings. However, the numberof stacking interactions is a few times larger than the number of CH. . .π interactions. Theanalysis also shows that in almost all cases, CH. . .π interactions are formed only whenstacking interactions are prevented by voluminous substituents. Hence, between planarchelate and C6-aryl rings stacking interactions are preferred to CH. . .π interactions.

10.3.2.3 Stacking Interactions between Chelate Rings

The data related to the intermolecular chelate. . .chelate stacking interactions in the crystalstructures obtained from the CSD were analysed using the geometric parameters shown

Figure 10.25 Histogram showing the number of structures with specific interactions: (I)stacking interactions between chelate and C6-aryl rings, (II) stacking interactions of C6-arylwith aromatic rings fused with a chelate ring, (III) CH. . .π interactions between C6-aryl-H andπ -system of chelate rings, (IV) other CH. . .π interactions, (V) CH. . .X interactions, and (VI)structures without any of these interactions.

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(a) (b)

Ct

Mʹ

M P2

P1

β

Mʹ

D

M

τ

Ω

Ωʹ

Figure 10.26 Geometric parameters describing the interactions between two chelate rings ofsquare-planar complexes. The chelate ring can be five- or six-membered. (a) Ct is the distancebetween the centres of two chelate rings, P1 and P2 are the mean planes of the chelate rings,and β is the angle between the normal to the chelate ring and the vector that connects thecentres of the two chelate rings. (b) D is the distance between the two metals, and τ is thetorsion angle metal-centroid-centroid0-metal0 (M-Ω-Ω‘-M’).

in Figure 10.26 [28]. It was considered that an interaction exists if ϕ < 10◦, Ct < 4.6A, and β < 35◦. To characterise the stacking interactions, the normal distances (R) fromthe centroid of one chelate ring to the mean plane of another and the torsion angles τ

(Figure 10.26) were also analysed.The analysis showed that chelate. . .chelate stacking interactions occur in a large number

of crystal structures of neutral square-planar complexes. Only square-planar complexeswithout charge were analysed in order to avoid the influence of the charges on the in-teractions. Chelate. . .chelate stacking interactions were found in 955 structures of neutralcomplexes. In a large number of these structures, there are two or more interactions. Hence,in these structures there are 1866 chelate–chelate stacking interactions (Table 10.7). Theseinteractions exist in the structures of complexes with various metal atoms and ligands, andthey are ubiquitous in structures of square-planar complexes. Many of the complexes havemore than one chelate ring in the coordination sphere and these chelate rings are in most

Table 10.7 Number of interactions in structures of neutral square-planar complexescontaining different types of chelate rings

Interactionsin neutralcomplexes Total Parallel Cross Antiparallel

five-membered 46 15 3 23six-membered 42 4 0 38

Interactionsof isolatedchelates

89 five- and six-membered

1 1 0 0

overall 89 20 3 61five-membered 1229 139 201 655six-membered 410 118 53 176

Interactionsof fusedchelates

1777 five- and six-membered

138 21 78 2

overall 1777 278 332 833

P1: TIX/XYZ P2: ABC



cases fused. Since it was shown that the size of the planar system has an influence on thestacking interactions, crystal structures of complexes with isolated and fused chelate ringswere analysed separately.

10.3.2.3.1 Stacking Interactions of Isolated Chelate Rings There are 151 crystal struc-tures of neutral square-planar complexes with isolated chelate rings where aromatic organicfragments were excluded to avoid their possible influence on the chelate–chelate interac-tions. The complexes had one or two isolated chelate rings. In 31 of the 88 structureswith five-membered chelate rings and in 32 of the 63 structures with six-membered rings,stacking interactions between chelate rings were observed. An analysis of the constitutionof five-membered chelate rings shows that these rings can be planar but, at the same time,have sp3-atoms in the ring – these chelate rings do not form stacking interactions.

Bulky substituents can make the stacking interaction of two chelate rings impossible.All the structures featuring six-membered chelate rings possess β-unsaturated ligands asanalogues or derivatives of acac complexes. The inspected structures with six-memberedrings show that the chelate. . .chelate interactions do not occur if unsaturated groups suchas CN, OCCH3, NCCH3 and OCOEt are substituents on the β-carbon atom of the chelatering. It seems that these groups disrupt the delocalisation within the chelate ring and renderthem incapable of forming stacking interactions.

According to the defined criteria, in many structures more than one interaction wasobserved. A further analysis of the structures with isolated chelate rings revealed 46interactions between five-membered, 42 interactions between six-membered, and oneinteraction between five- and six-membered isolated chelate rings (Table 10.7). Thereare five types of metal atoms represented and 29 types of chelate rings involved inchelate–chelate stacking interactions.

In order to show that close contacts of two molecules are the consequence of attrac-tive chelate. . .chelate interactions and not a consequence of other attractive interactionsbetween them, all interactions of the isolated chelate rings were manually evaluated. Nohydrogen bonds between two interacting molecules were found. In four structures, shortmetal–metal distances were observed, in three, CH–O interactions, and in one structure, aCH. . .π interaction was noted. Because of the small number of the structures with addi-tional interactions as well as the fact that the additional interactions were not consideredstrong, one can consider that attractive chelate. . .chelate interactions are responsible forthe close contacts.

The distribution of the torsion angles (τ ), that describe the mutual orientation of twochelate rings (Figure 10.26), shows that in most cases τ is less than 10◦ or near 180◦,that is, indicating that parallel and antiparallel conformations (Figure 10.27), respectively,dominate. For interactions involving five-membered chelate rings, the number of paralleland antiparallel orientations is similar. However, in the case of six-membered chelate rings,there are only a few examples having a parallel orientation, with most examples havingantiparallel orientations (Table 10.7). Data for five-membered rings show that there are afew cases of structures with a cross conformation (τ ∼90◦) (Figure 10.27).

The larger values of the β angle for five-membered chelate rings show that these com-plexes have a greater tendency for slipped-parallel orientations (Figure 10.28), which isin accordance with the most stable conformation of two benzene rings and in stackinginteractions of chelate rings with organic aromatic rings (Chapter 10.3.2.1). On the other

P1: TIX/XYZ P2: ABC



M

M M

M M

M

180°90° or 100°0°

Figure 10.27 Three most frequent conformations in crystal structures with chelate–chelatestacking interactions and values of torsion angle τ (Figure 10.26).

hand, for six-membered rings, values of the β angle are quite small, with most of thembelow 12◦, indicating that the two rings are in a face-to-face orientation with a very smalloffset. In order to enable stacking in spite of the substituents on the chelate rings, the inter-actions have an antiparallel conformation and small values of angle β to avoid clashing ofthe substituents (Figure 10.29).

10.3.2.3.2 Stacking Interactions of Fused Chelate Rings A search of the CSD revealed2322 square-planar complexes in which chelate rings are fused with other chelate or/andorganic aromatic rings, that is, with the chelate ring being part of an extended π -system.In 892 of these structures, stacking interactions were observed between two chelate rings.Hence, in the majority of these structures chelate rings do not form mutual stacking inter-actions. The main reason for this is the large size of the planar ligand. In complexes withlarge planar ligands (for example, bipy, phen and terpy), it is probable that aromatic organicfragments overlap mutually or with chelate rings, while two chelate rings do not over-lap mutually. This is shown in the analysis of the stacking interactions of heteroaromaticligands, that is, of terpy and phen complexes (Figure 10.30) [38].

In the 892 structures featuring chelate. . .chelate stacking interactions, 1777 stackinginteractions were observed between two chelate rings of which 1229 interactions occurredbetween five-membered chelate rings, 410 interactions between six-membered rings, and138 interactions between five- and six-membered chelate rings. These numbers indicatethat in most of the structures there is more than one interaction, which could be anticipated

Figure 10.28 Two views of the stacking interactions between bis(thiosemicarbazido)nickel(II)molecules in the structure of bis(cis-bis(thiosemicarbazido)nickel(II)) trans-diaqua-bis(malonato)nickel(II) dinitrate tetrahydrate. In this structure there are three chelate–chelatestacking interactions; two of them are in parallel conformation (showed with dashed lines),and one is antiparallel (showed with solid line).

P1: TIX/XYZ P2: ABC



Figure 10.29 Two views of the six-membered chelate–chelate stacking interaction in theantiparallel conformation (τ ∼180◦) in the structure of bis(acetylacetonato)copper(II).

since in most of the complexes there are two or three fused chelate rings. Similar to the datafor isolated chelate rings, for fused rings the torsion angle τ (Figure 10.26) has preferredvalues, being close to 180◦ in a substantial number of cases. However, there are fewerinteractions where this angle is less than 10◦ and even fewer with values around 90◦. Theseresults indicate that in crystal structures with fused chelate rings there are also three mainconformations of interacting chelate rings; parallel, antiparallel and crossed (Figure 10.27).The distributions of normal distances of two interacting chelate rings show that the normaldistances are in the range typical for stacking interactions found for all-organic systems.

The data indicate somewhat shorter normal distances for six-membered rings. The peak ofthe distributions for the interactions of six-membered rings is in the range 3.3–3.5 A. For theinteractions of five-membered rings, the peak is in the range 3.4–3.5 A with large numbersof interactions with R above 3.5 A. While fused six-membered rings show a tendency forsomewhat smaller values of the β angle, the difference between five- and six-memberedrings is much less pronounced than in the case of isolated five- and six-membered chelaterings. The values of the β angle for fused chelate rings are in a large range, indicatingslipped-parallel orientations (values of angle β above 15◦) but also almost face-to-faceorientations (small values of angle β), that are not typical for stacking interactions betweenbenzene molecules and between chelate rings with C6 aromatic rings (Section 10.3.2.1).The large range of values of the β angle for fused chelate rings is a consequence of thestacking of a whole fused, large planar system.

Figure 10.30 Two projections presenting stacking interaction in crystal struc-ture (1R, 2R)-1,2-diaminocyclohexane-N,N’)(1,10-phenanthroline-N,N’)-platinum(II) bis(hexafluorophosphate). In this structure there are no chelate–chelate stacking interactions.

P1: TIX/XYZ P2: ABC



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