1.2. Noncovalent interactions
The weak noncovalent interactions have a constitutive role in biological or biomimetic systems as well as in artificial supramolecular structures. Noncovalent or van der Waals interactions were first recognized by J. D. van der Waals in the nineteenth century [3]. Their role in nature has been unravelled only during the past two decades.
In contrast to the covalent interactions that dominate in classical molecules, noncovalent interactions are weak interactions that bind together different kinds of building blocks into supramolecular entities [3,4]. Covalent bonds are generally shorter than 2 Å, while noncovalent interactions function within range of several angstoms. The formation of a covalent bond require overlapping of partially occupied orbitals of interacting atoms, which share a pair of electrons. In noncovalent interactions, in turn, no overlapping is necessary because the attraction comes from the electrical properties of the building blocks.
The noncovalent interactions or van der Waals forces involved in supramolecular entities may be a combination of several interactions, e.g. ion-pairing, hydrogen bonding, cation−π , π −π interactions etc. [1-5]. A wide range of attractive and repulsive forces is subsumed under term noncovalent. Noncovalent interactions comprise interactions between permanent multipoles, between a permanent multipole and an induced multipole, and between a time–variable multipole and an induced multipole. The stabilizing energy of noncovalent complexes is generally said to consist of the following energy contributions: electrostatic (or Coulombic), induction, charge transfer, and dispersion. These terms are basically attractive terms. The repulsive contribution, which is called exchange-repulsion, prevents the subsystems from drawing too close together. The term induction refers to general ability of charged molecules to polarize neigbouring species, and dispersion (London) interaction results from the interactions between fluctuating multipoles. In charge-transfer (CT) interactions the electron flow from the donor to the acceptor is indicated. The term van der Waals (vdW) forces is frequently used to describe dispersion and exchange-repulsion contributions, but sometimes also other long–range contributions are included in the definition. All of these interactions involve host and guest as well as their surroundings (e.g. solvation, crystal lattice and gas phase).
1.2.1. Ion pairing
In ion pairing (ion–ion, ion–dipole, dipole–dipole etc.) the driving force is electrostatic (Coulombic) interactions, which unquestionably play an important role in natural and in supramolecular systems [1,4]. Charges are heavily delocalized in organic ions, which complicates the theoretical analysis of ion pairing [4,5,9]. As a means to understand ion pairs, theoretical approaches to the association constant (K) based on Debye-Hückel theory have been developed by Bjerrum (spherical ions with point charges) and Fuoss (contact ion pairs). A numerical method introduced by Poisson permits the consideration of solvent molecules, and the salt effect is well described by Manning’s counterion condensation theory [4,9]. As an example of supramolecular ion–ion interaction can be presented the interaction of organic cation tris(diazabicyclooctane) with Fe(CN)63- [1]. The structure of alkali metal cation with macrocyclic crown ether can be presented as an example of supramolecular ion–dipole interaction. In this structure the oxygen lone electron pairs are attracted to the cation positive charge. Between neutral polar molecules the electrostatic contributions comes mostly from dipole–dipole interactions.
1.2.2. Hydrogen bonding
Hydrogen bonding is a relatively strong and probably the most important noncovalent interaction [1,3-5]. Hydrogen bonded complexes are stabilized by electrostatic, induction and dispersion energy terms. The electrostatic term is contibuted by dipole–dipole and ion–dipole interactions, which give hydrogen bonds their highly directional nature. A hydrogen bond, D–H···A, is formed between a hydrogen attached electronegative donor atom (D) and a neigbouring acceptor atom (A). Different types of H-bonds are presented in Fig. 1, where the two last types are the less frequent ones [4]. Properties and a common classification of the H-bonds in terms of strength are represented in Table 1. The concept of hydrogen bonding has also been extended to the weaker C–H···O type, which has been studied systematically only recently [10-13].
Table 1. Classification and some properties of hydrogen bonds. [1]
| D-H···A interaction | Strong | Moderate | Weak |
|---|---|---|---|
| Mainly covalent | Mainly electrostatic | Electrostatic | |
| Bond energy ( kJmol -1) | 60–120 | 16–60 | < 12 |
| Bond lengths (Å): H···A | 1.2–1.5 | 1.5–2.2 | 2.2–3.2 |
| D···A | 2.2–2.5 | 2.5–3.2 | 3.2–4.0 |
| Bond angle ( º ) | 175–180 | 130–180 | 90–150 |
| Examples | Gas phase dimers with strong acids/bases, HF complexes | Acids, Biological molecules | C-H···N/O and N/O–H···p hydrogen bonds |
1.2.3. Cation–π interaction
In cation−π interactions surprisingly strong forces are found between cations and a π -face of an aromatic structure [1,3,4,14]. Electrostatic forces appear to play the dominant role in the cation−π interaction, though modern theories also involve additional terms such as induced dipole, polarizability, dispersion and CT. The role of electrostatic interactions and the charge distribution of aromatics are illustrated in Fig. 2b, where the quadrupole moment of benzene is shown in emphasized form. Cation−π interactions are well described by a schematic drawing such as Fig. 2a showing a K+ ion interacting with the negatively charged π -electron cloud of benzene.
1.2.4. Aromatic π −π interaction
Weak electrostatic, π −π interactions occur between aromatic moieties [1,4,15,16]. The stabilizing energy of π −π interactions also includes induced dipole and dispersion contributions. Two general types of aromatic π −π interactions are face-to-face and edge-to-face (Fig. 3a). The latter is actually a C−H···π interaction (the C−H bond generally having a small dipole moment). The attraction in these two orientations comes from the interaction between positively charged hydrogen atoms and negatively charged π -face of aromatic system. The perfect facial alignment of face-to-face orientation is unlikely because of the electrostatic repulsion between the two negatively charged π -systems of aromatic rings (Fig. 3b). The distance between the aromatic π −π faces is about 3.3−3.8 Å.
