Nanosystems Intermolecular Forces and Potentials

"It will be perfectly clear that in all my studies I was quite convinced of the real existence of molecules, that I never regarded them as a figment of my imagination, nor even as mere centres of force effects. " J. D. van der Waals


In this chapter, we present details of interatomic and intermolecular forces and potential energy functions from the point of view of formation and functions of nanostructures. Detailed information about interatomic and intermolecular interactions is necessary for modeling, prediction and simulation of the behavior of assembly of finite number of molecules, which happen to exist in nanosystems. In addition such information will guide the design of positional-assembly and prediction of self-assembly and self-replication which are the fundamentals of bottom-up nanotechnology.

In this chapter, we first introduce a short description of covalent interactions and their differences with non-covalent interactions. Then we will present a detailed analysis of the non-covalent and covalent interactions. This will include experimental and theoretical modeling followed by phenomenological models developed for non-covalent and covalent interactions.

Covalent and Noncovalent Interactions

The interactions between atoms and molecules are either of covalent or non-covalent type [1,2]. Covalent interactions are in a form of chemical bond in compounds that result from the sharing of one or more pairs of electrons. Atoms can combine by forming molecules to achieve an octet of valence electrons by sharing electrons. In this process, the electron clouds of every two atoms overlap where they are thicker and their electric charge is stronger. Since the nuclei of the two atoms have stronger attractive potentials to their respective thick electron clouds than their mutual repulsive potential (due to their far distance from one another), the two atoms are held together and form a molecule (see Figure 1).

Figure 1. Symbolic demonstration of covalent bond formation between two atoms.

As a result, the formation of covalent bonds requires overlapping of partially occupied orbitals of interacting atoms, which share a pair of electrons. On the other hand, in non-covalent interactions no overlapping is necessary because the attraction comes from the electrical properties of the interacting atoms and molecules. The electrons in the outermost shell are named the valence electrons, which are the electrons on an atom that can be gained or lost in a chemical reaction. The number of covalent bonds, which an atom can form, depends on how many valence electrons it has. Generally, covalent bond implies the sharing of just a single pair of electrons. The sharing of two pairs of electrons is called a double bond and three pairs sharing is called a triple bond. The latter case is relatively rare in nature, and actually two atoms are not observed to bond more than triply. Most frequently covalent bonding occurs between atoms, which possess similar electronegativities. That is when they have similar affinities for attracting electrons. This occurs when neither atom can possess sufficient affinity, or attractive potential energy, for an electron to completely remove an electron from the other atom.

Covalent interactions are stronger than the non-covalent bonds. The weak non-covalent interactions were first recognized by J. D. van der Waals in the nineteenth century [2, 3]. Covalent interactions are of short range and the resulting bonds are generally less than 0.2 0 [nm] long. The non-covalent interactions are of longer range and within a few 0 [nm] range. In Table 1 the differences between covalent and non-covalent interactions are reported.

Table 1. Differences between covalent and non-covalent interactions.






van der Waals, hydrogen bonding, etc.

ENERGETIC (kcal/mol)





Open to change


Mainly AH

Both AH and TAS


Chemical reaction

Chemical equilibrium





Less important

Very important

0 0

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