Why is bond length affected by multiple bonding?

Bond length is affected by multiple bonding because multiple bonds are shorter and stronger than single bonds.

In more detail, the concept of multiple bonding refers to the sharing of more than one pair of electrons between two atoms. This can occur in the form of double or triple bonds. The presence of multiple bonds significantly influences the bond length, which is the distance between the nuclei of two bonded atoms.

The reason behind this is rooted in the nature of the bonding itself. In a single bond, one pair of electrons is shared between the atoms. However, in a double bond, two pairs of electrons are shared, and in a triple bond, three pairs are shared. This increased sharing of electrons results in a stronger electrostatic attraction between the positively charged nuclei and the negatively charged electrons. This stronger attraction pulls the nuclei closer together, resulting in a shorter bond length.

For instance, consider the carbon-carbon bonds in ethane (C2H6), ethene (C2H4), and ethyne (C2H2). Ethane has a single bond, ethene has a double bond, and ethyne has a triple bond. The bond length decreases from ethane to ethyne, with the bond in ethyne being the shortest. This is because the triple bond in ethyne involves more shared electron pairs, creating a stronger bond and consequently a shorter bond length.

Moreover, the type of atomic orbitals involved in the bond formation also plays a role. In multiple bonds, there is involvement of both sigma (σ) and pi (π) bonds. A sigma bond is formed by the end-to-end overlapping of atomic orbitals, while a pi bond is formed by the side-to-side overlapping. Sigma bonds are stronger and shorter than pi bonds. However, the presence of a pi bond in addition to a sigma bond in multiple bonds still results in an overall stronger and shorter bond compared to a single bond.

In conclusion, multiple bonding affects bond length by creating stronger electrostatic attractions between the bonded atoms, leading to shorter and stronger bonds. This is influenced by the number of shared electron pairs and the type of atomic orbitals involved in the bond formation.