Final answer:
Covalent bonds formed with atoms that have smaller atomic radii tend to be stronger because the bonding electrons are closer to the nucleus. Thus, in the case of CX versus CY, if X and Y have the same electronegativity but Y has a smaller radius, CY would have a higher bond enthalpy.
Step-by-step explanation:
The strength of a covalent bond is influenced by the atomic radii of the atoms involved. A smaller atomic radius means the bonding electrons are closer to the nucleus, leading to a stronger attractive force between the nucleus and electron clouds. Therefore, in a molecule of CX compared to CY, if nonmetal X has a larger atomic radius than nonmetal Y and they have equivalent electronegativities, the bond in CY would generally have a higher bond enthalpy. This is because the shorter distance between the nucleus and the bonding electrons in Y allows for a stronger covalent bond compared to X.
In other words, the smaller the covalent radius, the closer the bonding electrons are held to the nucleus. As a result, the bonds formed with smaller radii are stronger compared to those formed with atoms having larger radii, assuming other factors like electronegativity are equal.
The strength of a covalent bond is influenced by the atomic radii of the atoms involved. In general, the bond strength increases as the atomic radius decreases. This is because when the atomic radii are smaller, the shared electrons are closer to the nuclei of the atoms, resulting in stronger electrostatic forces of attraction between the nuclei and the electrons.
In the scenario you presented, if the bond enthalpy is defined as the amount of energy required to break a bond, then the molecule with a higher bond enthalpy would be CX. This is because carbon has a smaller atomic radius compared to nonmetal atom X, so the shared electrons will be closer to the carbon nucleus, resulting in a stronger covalent bond.