We first argue that the covalent bond and the various closed-shell interactions can be thought of as symmetry broken versions of one and the same interaction, viz., the multi-center bond. We use specially chosen molecular units to show that the symmetry breaking is controlled by density and electronegativity variation. We show that the bond order changes with bond deformation but in a step-like fashion, regions of near constancy separated by electronic localization transitions. These will often cause displacive transitions as well so that the bond strength, order, and length are established self-consistently. We further argue on the inherent relation of the covalent, closed-shell, and multi-center interactions with ionic and metallic bonding. All of these interactions can be viewed as distinct sectors on a phase diagram with density and electronegativity variation as control variables; the ionic and covalent/secondary sectors are associated with on-site and bond-order charge density wave respectively, the metallic sector with an electronic fluid. While displaying a contiguity at low densities, the metallic and ionic interactions represent distinct phases separated by discontinuous transitions at sufficiently high densities. Multi-center interactions emerge as a hybrid of the metallic and ionic bond that results from spatial coexistence of delocalized and localized electrons. In the present description, the issue of the stability of a compound is that of mutual miscibility of electronic fluids with distinct degrees of electron localization, supra-atomic ordering in complex inorganic compounds comes about naturally. The notions of electronic localization advanced hereby suggest a high throughput, automated procedure for screening candidate compounds and structures with regard to stability, without the need for computationally costly geometric optimization.
I. MOTIVATIONChemical bonding is traditionally discussed in terms of the covalent, ionic, and metallic bond 1 , and weaker, closed-shell interactions such as secondary, donoracceptor, hydrogen, and van der Waals.2,3 The distinction between these canonical bond types is not always clear-cut. For instance, a directional, multi-center 4 bond holding together identical atoms has an inherent ionic feature: In a three-center, linear ppσ bond, 5,6 the central atom contributes only a half orbital to each of the individual ppσ bonds.7 The terminal atoms, on the other hand, each contribute one full orbital, implying a non-uniform charge distribution over the bond. At the same time, the three-center ppσ bond can be thought of as a limiting case of the metallic bond, since the appropriate electron count for an infinite chain corresponds to a half-filled band.4 This identification is consistent with the metallic luster of compounds in which covalent and secondary bonds are comparable in length.3 . In solid-state contexts, interplay between ionic and covalent interactions is often discussed using the van Arkel-Ketelaar triangle for binary compounds;8-11 or revealed, for insta...