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This chapter presents the results obtained when the level of student preparedness for 2nd year Inorganic chemistry module was investigated. This was done by using a diagnostic test based mainly on concepts that are pre-requisite or foundational knowledge (high school and 1st year chemistry syllabi) to the 2nd year Inorganic Chemistry course, CHE 2521. The diagnostic test covered the following concepts: hybridization, electronegativity, octet rule, bonding, periodic table, electronic configuration, Lewis structures and resonance. The study highlighted the presence of misconceptions and general lack of preparedness for the subsequent CHE 2521 module. The most important misconceptions suggested by students’ responses are described and discussed, and, where possible, suggestions on their sources are outlined.

Though almost a century old, Lewis’s theory of chemical bonding remains at the heart of the understanding of chemical structure. In spite of their basic discrete nature, Lewis’s structures (topological 0-manifolds) continue to lend themselves to sophisticated treatments leading to valuable results in terms of topological analysis of chemical properties. The bonding topology is however not only defined, but also refined by direct consideration of the nuclear geometry, itself determined by the configuration of the embedding electron cloud. During the last century, the theory has thus been complemented by the mesomery concept, by the Linnett’s double quartet scheme and by the VSEPR/LCP models. These models rely on an assumed spatial disposition of the electrons which does not take the quantum mechanical aspects into account. These models are reexamined by investigation of the topological 1-manifolds generated by the gradient field of potential functions featuring the electron cloud configuration, such as the electron density or electron localization function (ELF). In this chapter, we reexamine these models in order to escape from the quantum mechanical dilemma and we show how topological analyzes enable to recover these models.

The seminal papers of Lewis and Kossel in 1916 are put into a historical perspective. Mendeleev’s periodic table, Thompson’s discovery of the electron, Ramsay and Raleigh’s discovery of the noble gases, Rutherford’s model of the atom and Bohr’s description of the stationary orbitals for the electrons in atoms all paid an important role in providing the background for Lewis and Kossel’s proposal that the chemical bond originated either from the transfer of electrons or the sharing of electron pairs. These insights depended on the attainment of inert gas configurations by the atoms either directly by electron transfer or electron-pair sharing. The model incorporated an evolutionary gene which has enabled it to survive and grow by incorporating subsequent developments in quantum physics. The simplicity of the model has resulted in the development of a notation, which is universally used by chemists and has evolved to plot the course of chemical reactions and predict their regioselectivities. Its initial limitations are discussed, and the way in which they have been overcome by an orbitally based model is recounted. The model has been repeatedly enriched by quantum mechanically based theoretical studies.

The electron-pair bonding model that was introduced by Gilbert Lewis 100 years ago is discussed in the light of modern quantum chemical methods for analysing the electronic structures of some simple molecules. It is argued that Lewis structures in conjunction with accurate quantum chemical calculations are still very useful for the description of chemical bonding. The emphasis lies on the difference between electron-sharing bonds A–B and dative bonds A → B which were suggested by Lewis as a general definition for acids and bases. The electron-pair model, if combined with quantum chemical calculations, remains a powerful guide for the search of new molecules and for understanding molecular structures.

In elementary chemical bonding theory, a rather neglected concept is a type of chemical bond that Pauling has designated as the “3-electron bond”. This type of bond is usually represented as A···B and it involves three electrons distributed amongst two overlapping atomic orbitals centred on the atoms A and B. Alternative designations are either “3-electron 2-centre” bond or “3-electron 2-orbital” bond.

Most of the “increased-valence” structures that we have discussed so far may be derived from Lewis structures by delocalizing lone-pair \( \uppi \) and/or \( {\overline{\uppi }} \) electrons into vacant bonding or antibonding orbitals. The atomic orbital overlaps that are appropriate for some of these delocalizations are shown in Figs. 1-5 and 2-4. We shall now consider a few systems whose “increased-valence” structures can be constructed by delocalizing one or more lone-pair σ electrons of a Lewis structure into bonding or antibonding σ orbitals. Some other examples will also be discussed in Chapter 20, where the theory will be presented in a slightly different form. However, the principles for both chapters are the same.

The technique described in Chapter 12 for constructing “increased-valence” structures, namely that of delocalizing one or more non-bonding electrons of a standard Lewis structure into adjacent bonding orbitals, is quite general and easily applied. We shall now use this method to construct “increased-valence” structures for numerous molecular systems that involve 6-electron 5-centre, 8-electron 6-centre and longer N-centre bonding units, as well as for some molecules with 4-electron 3- centre and 6-electron 4-centre bonding units. In general, we shall find that only one or two “increased-valence” structures are required in order to make deductions concerning bond lengths that are in qualitative accord with the measured lengths, i.e. for the systems considered, “increased-valence” structures provide easily derived and economical representations of their electronic structures.

We are now ready to examine in detail the incorporation of the Pauling “3-electron bond” structure Ȧ · Ḃ into the valence-bond structures for electron-rich molecules that involve 4-electron 3-centre and 6-electron 4-centre bonding units. To do this, we may use any of three alternative methods. In this Chapter, we shall discuss one of them. It involves the spin-pairing of the unpaired-electron of Ȧ · Ḃ with the unpaired electron of either an atom Ẏ or a second Pauling “3-electron bond” structure Ċ · Ḋ.

Mulliken1–3 has provided a quantum mechanical description of molecular complexes that are formed by reaction between an electron-donor (D) and an electron acceptor (A). The wave-function for the normal or ground-state of the complex may be expressed1–3 approximately according to Eqn. (1).

Valence-bond structures with Pauling “3-electron bonds” between pairs of atoms may be written down for a number of triatomic radicals. Here we shall examine these types of structures for some radicals with either 17 or 19 valence-shell electrons. For these systems, it is necessary to write down two Pauling “3-electron bond” structures, which participate in resonance. The delocalized molecular orbital equivalent of this resonance involves the construction of three 3-centre molecular orbitals to accommodate five electrons; this is described in Section 6-4.

Transition metal complexes have been playing an increasingly important role in modern chemistry in the past century, and this is partly due to their distinctive structure and bonding features that allow them to play a special role in organometallic reactions. Despite their importance, the current understanding of their structure and bonding relies to a large extent on sophisticated quantum chemical treatments, which do not encourage the formulation of more generalized rules. In this review, commemorating the centennial anniversary of the seminal Lewis paper, we would like to go back to basics and start from the classical Lewis description and then combine some observations we obtain from modern molecular orbital theory to give a simple but general bonding picture for transition metal complexes. This model, albeit simple, provides a localized description to metal–ligand interactions in these complexes and allows us to easily treat various cases with atypical metal–ligand or even metal–metal interactions in a modular manner.