Generally, as we move down the halogen group, the atomic size increases, which leads to longer and weaker bonds, so bond dissociation enthalpy should decrease . However, fluorine ($\text{F}_2$) is anomalous. Because of the extremely small size of the fluorine atom, the lone pairs of electrons on the two adjacent bonded atoms experience strong interelectronic repulsions , making the $\text{F-F}$ bond weaker than expected. As a result, the bond dissociation enthalpy of $\text{F}_2$ is lower than that of both $\text{Cl}_2$ and $\text{Br}_2$. According to the standard values , the bond dissociation enthalpies are roughly $\text{Cl}_2$ ($\approx 243$ kJ/mol) > $\text{Br}_2$ ($\approx 192$ kJ/mol) > $\text{F}_2$ ($\approx 155$ kJ/mol) > $\text{I}_2$ ($\approx 151$ kJ/mol). Therefore, the correct order is $\text{Cl}_2 > \text{Br}_2 > \text{F}_2 > \text{I}_2$.