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2 Rational Design of Superatoms Using Electron‐Counting Rules

       Puru Jena1, Hong Fang1, and Qiang Sun2,3

       1 Physics Department, Virginia Commonwealth University, Richmond, Virginia, USA

       2 School of Materials Science and Engineering, Peking University, Beijing, China

       3 Center for Applied Physics and Technology, Peking University, Beijing, China

2.1 Introduction

The periodic table of elements developed by Mendeleev and presented at a meeting of the Russian Chemical Society on 6 March 1869, was based purely on the chemical properties of the then known elements. However, a fundamental understanding of the chemistry of the elements and the sites they occupy in the periodic table had to wait until the discovery of the electron in 1897 and the development of quantum mechanics in the early part of the twentieth century. Central to this understanding is the knowledge of the atomic orbitals and the manner in which they are filled as one moves along the columns and the rows of the periodic table. The elements belonging to the same column of the periodic table have similar chemistry. As one moves along the rows, the electrons continue to fill the successive atomic orbitals in keeping with the Pauli's principle and the chemistry changes accordingly. Consider, for example, the Group 18 elements, which have an outer electron configuration of ns ² np ⁶. Because these orbitals are full, the energy cost to remove an electron is high and the energy gain in adding an electron is negligibly small. Thus, these elements do not participate in chemical reactions, justifying their name as the noble gas atoms. They are very stable and their bonding is characterized by weak van der Waals forces. The origin of their stability is, therefore, attributed to the electronic shell closure, which is referred to as the “octet rule” where 2 + 6 = 8 electrons are enough to fill the s and p orbitals. On the other hand, the group 1 alkali elements are characterized by an outer electron orbital occupation, namely ns ¹. The energy cost to remove this electron, i.e., the ionization potential, is smaller than elements belonging to the same row in the periodic table and hence in a chemical reaction they tend to donate that electron, leaving behind an ionic core with closed electronic shells. Similarly, group 17 halogen atoms have an outer electron configuration of ns ² np ⁵ and need an extra electron to satisfy the octet rule. Consequently, they have large electron affinities and gain energy by accepting an electron in a chemical reaction. Thus, when an alkali atom and a halogen atom approach each other, electron transfer from the alkali to the halogen atom satisfies the octet rule of both the atoms, resulting in the formation of a salt and an ionic bond between the cations and anions. In other cases, such as H₂, O₂, and N₂, electron shell closure is achieved not by transferring electrons from one atom to the other but rather by sharing electrons. This leads to a covalent bond, which, in general, is stronger than an ionic bond. A fourth kind of bonding appears as atoms from groups 2–13 and some elements in the higher groups upto 16 come together to form a crystal. Here, each atom contributes its outer valence electrons to a common pool. These electrons move “freely” and collectively, forming a metal. Note that, the ionic cores of the metal also have electronic shell closure.

The periodic table (see Figure 2.1) currently consists of 118 elements among which 94 occur in nature. All materials are created by combining atoms from one or more groups. Among the 94 naturally occurring elements, some are expensive while some others occur in trace quantities. Some of these elements are even toxic. Is it possible to replace the expensive elements by earth abundant materials and the toxic elements by nontoxic ones? This has been the dream of alchemists for centuries. With the birth of cluster science, we have arrived at a stage where this dream may not be as farfetched as it once seemed.

Atomic clusters are groups of atoms whose size and composition can be varied by design, one atom at a time. More than half a century of research has made it clear that the properties of clusters are very different from any other form of matter [1]. Because their properties are size‐, composition‐, and shape‐specific, clusters can be tailored with atomic precision. In 1992, Khanna and Jena [2] coined the word “superatom” to describe a cluster that has the same chemistry as an atom in the periodic table and suggested that these superatoms can be used as the building blocks of a new three‐dimensional periodic table, with superatoms forming the third dimension [3]. If these superatoms can retain their geometry and properties when assembled, a new class of cluster‐assembled materials with tailored properties can be formed. A classic example of such a crystal is based on C₆₀ fullerene, which was discovered by Smalley and coworkers in the gas phase in 1985 [4] and later synthesized in bulk quantities by Kratschmer et al. [5]. Once assembled, C₆₀ fullerenes retain their shape, but the property of the fulleride crystal is very different from that of graphite and diamond. The former is the ground state of carbon while the latter is metastable but protected by a very large energy barrier. Note that the discovery of C₆₀ was not the result of a rational design approach. The question is: can other clusters like C₆₀ be rationally designed by using some prescribed rules?