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In coordination chemistry, a coordinate covalent bond, [1] also known as a dative bond, [2] dipolar bond, [1] or coordinate bond [3] is a kind of two-center, two-electron covalent bond in which the two electrons derive from the same atom. The bonding of metal ions to ligands involves this kind of interaction. [4]
Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. [18] When each atom contributed one electron to the bond, it was called a covalent bond. When both electrons come from one of the atoms, it was called a dative covalent bond or coordinate bond. The distinction is not very clear-cut.
For doubly bridging (μ 2-) ligands, two limiting representation are 4-electron and 2-electron bonding interactions. These cases are illustrated in main group chemistry by [Me 2 Al(μ 2-Cl)] 2 and [Me 2 Al(μ 2-Me)] 2. Complicating this analysis is the possibility of metal–metal bonding.
The covalent bond classification (CBC) method, also referred to as LXZ notation, is a way of describing covalent compounds such as organometallic complexes in a way that is not prone to limitations resulting from the definition of oxidation state. [1]
A number of organic compounds form charge-transfer complex, which are often described as electron-donor-acceptor complexes (EDA complexes). Typical acceptors are nitrobenzenes or tetracyanoethylene (TCNE). The strength of their interaction with electron donors correlates with the ionization potentials of the components.
Oxidation of R 3 P–M complexes results in longer M–P bonds and shorter P–C bonds, consistent with π-backbonding. [11] In early work, phosphine ligands were thought to utilize 3d orbitals to form M–P pi-bonding, but it is now accepted that d-orbitals on phosphorus are not involved in bonding as they are too high in energy. [12] [13]
In Organic chemistry, the inductive effect in a molecule is a local change in the electron density due to electron-withdrawing or electron-donating groups elsewhere in the molecule, resulting in a permanent dipole in a bond. [1] It is present in a σ (sigma) bond, unlike the electromeric effect which is present in a π (pi) bond.
Formal charges in ozone and the nitrate anion. In chemistry, a formal charge (F.C. or q*), in the covalent view of chemical bonding, is the hypothetical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.