What is the Difference Between dπ-dπ Bond and Delta Bond?

The key difference between a dπ-dπ bond and a delta bond lies in the atomic orbitals involved in their formation and the number of nodal planes present in the bond.

• dπ-dπ bond: This type of covalent chemical bond forms between a filled d atomic orbital and an empty d atomic orbital. The overlapping of these orbitals creates a coordinate bond. dπ-dπ bonds are typically observed in coordination compounds, where a metal binds with a ligand through the sharing of electron pairs.
• Delta bond: A delta bond is a covalent chemical bond formed by the overlapping of two d orbitals. Delta bonds have two nodal planes that go through both atoms involved in the bond. These bonds are usually observed in organometallic species, such as some ruthenium and molybdenum compounds.

In summary, a dπ-dπ bond forms between a filled and an empty d atomic orbital, while a delta bond is formed by the overlapping of two d orbitals. The number of nodal planes in a dπ-dπ bond is one, whereas a delta bond has two nodal planes.

Comparative Table: dπ-dπ Bond vs Delta Bond

Here is a table comparing the differences between dπ-dπ bonds and delta bonds:

In summary, dπ-dπ bonds form between a filled and an empty d atomic orbital, typically in coordination complexes involving transition metals and ligands. Delta bonds, on the other hand, are less common and involve face-to-face overlap of d-orbitals. They are generally weaker than σ and π bonds and have two nodal planes along the bond.