The extent of overlap between atomic orbitals generally decreases as you move down a group (column) in the periodic table. This is because the atomic orbitals become larger as you go down a group, so they are less likely to overlap with each other.

In the series F-Cl-Br-I2, the extent of overlap between the atomic orbitals would be expected to decrease in the following order:

$$F > Cl > Br>I_2$$

This is because the atomic orbitals of fluorine are smaller than those of chlorine, which are in turn smaller than those of bromine, and so on. The smaller the atomic orbitals, the more likely they are to overlap with each other.

There are some exceptions to this general rule. For example, the extent of overlap between the 2p orbitals of oxygen and fluorine is actually greater than the extent of overlap between the 2p orbitals of nitrogen and fluorine. This is because the oxygen 2p orbitals are more diffuse than the nitrogen 2p orbitals, so they are more likely to overlap with the fluorine 2p orbitals.

The extent of overlap between atomic orbitals is an important factor in determining the strength of a chemical bond. The greater the extent of overlap, the stronger the bond. This is because the electrons in the overlapping orbitals are more likely to be shared between the atoms, which leads to a stronger attraction between the atoms.