In recent years, chalcogen atoms have been increasingly employed in organic molecules for their ability to develop an “atypical” type of non-covalent attractive interaction with nucleophiles, which are now termed chalcogen bonds (ChBs).¹  Despite their electron lone pairs, chalcogens also embody a Lewis acidic role, analogous to the one observed for main group elements from groups 14–15 and 17–18 in the periodic table. More generally, these interactions are specific cases of secondary bonding, that have a distance longer than “primary” covalent bonds, but shorter than the sum of the van der Waals radii (∑r_vdw) of the atoms involved.²  The origin of this interaction resides in the anisotropic distribution of the electron density around covalently bonded atoms, which generates a region of positive electrostatic potential collinear but opposite to a σ-bond, and thus called a σ-hole.³  This depletion of electrons leads to connections with areas with a higher electron density, such as lone-pair possessing atoms or anions. Since chalcogen atoms typically form two covalent bonds R–Ch, two σ-holes can be located opposite to these bonds (Figure 18.1a). Their magnitude is enhanced by increasing the electronegativity of the R substituents attached to Ch, either by increasing the polarizability of Ch itself.^4,5  However, several studies have proven that van der Waals dispersion forces and orbital mixing need to be considered^6,7  to fully rationalize the roots of ChBs. Specifically, a n²(A₂) → σ*(R₁–Ch) interaction takes place when R–Ch⋯A chalcogen bonds occur, where non-bonding electrons of electron-donating atom A (ChB acceptor) interact with the empty antibonding molecular orbital located on the chalcogen atom Ch (ChB donor, Figure 18.1b).