Q.Bond dissociation enthalpy of F$_2$ (158.8 KJ mol$^{-1}$) is lower than that of Cl$_2$ (242.6 KJ mol$^{-1}$) Why ?

[!NOTE]
The unit is printed "KJ mol$^{-1}$" (capital K) and the stem runs straight into "Why ?" with no punctuation before it -- both the book's own style, transcribed as printed.

In the tiny F$_2$ molecule the lone pairs on the two fluorine atoms sit so close together that they strongly repel each other and weaken the F - F bond; in Cl$_2$ the lone pairs are farther apart and the repulsion is negligible. …

Step 1. Fluorine has a smaller atomic size than chlorine, so in the F$_2$ molecule the two nuclei -- and with them each atom's three lone pairs -- are held very close together.

Step 2. The lone pairs on each F atom are so close that they strongly repel each other; this repulsion works against the bonding attraction and makes the F - F bond weak, so less energy is required to break it. The textbook illustrates this inside the Problem box (its "Fig. 7.1"):

Comparing lone pair - lone pair repulsion in the two halogen molecules: F2's small size forces its lone pairs close enough t …

• Assuming bond strength smoothly decreases down the group and so expecting F-F to be the strongest halogen-halogen bond -- lone-pair repulsion makes F2 anomalously weak (the real order is Cl-Cl > Br-Br > F-F > I-I). …