Final answer:
In chemical bonding, 'side-on' and 'end-on' refer to the orientation of orbital overlap forming sigma and pi bonds, respectively. Sigma bonds occur with end-to-end orbital overlap, concentrating electron density between the nuclei. Pi bonds form through side-by-side overlap, with electron density above and below the internuclear axis.
Step-by-step explanation:
When discussing chemical bonding in molecules, specifically in valence bond theory, the terms 'side-on' and 'end-on' refer to the orientation of orbital overlap that forms different types of covalent bonds: sigma bonds (σ bonds) and pi bonds (π bonds), respectively. A sigma bond is characterized by the end-to-end overlap of orbitals along the internuclear axis, which means that the electron density is most concentrated between the nuclei of the bonding atoms. This can occur in various bonding situations, such as the overlap of two s orbitals (H₂), the overlap of an s orbital with a p orbital (HCl), or the overlap of two p orbitals (Cl₂).
On the other hand, a pi bond is formed by the side-by-side overlap of p orbitals, resulting in electron density being concentrated above and below the internuclear axis, with a nodal plane that runs through the axis. In molecules like ethene (C₂H₄), the pi bond is represented by the second bond in the double bonds between the carbon atoms. For instance, in an ethyne (C₂H₂) molecule, the single bond is a sigma bond, while double and triple bonds consist of one sigma bond and one or two pi bonds respectively.
Understanding these concepts is crucial in chemical bonding theories and helps predict the behavior, shape, and strength of different molecules. In educational contexts, students may encounter questions asking them to identify sigma and pi bonds from illustrations of molecular orbitals, such as a Check Your Learning exercise where they have to categorize overlaps of orbitals as sigma or pi bonds.