Hi Mam,

For this question I was wondering why can't electronegativity differences also be a right answer when considering dipole movement?

ΔEN cannot be the correct answer here.
The electronegativity difference in S–Cl is actually less than in Be–Cl. So, based only on ΔEN, BeCl₂ should have the more polar bonds. That part is true.
However, bond polarity and molecular polarity are not the same thing.
In BeCl₂, each Be–Cl bond is polar, but the molecule is linear:
Cl — Be — Cl
Because the two identical bond dipoles are equal and point in opposite directions, they cancel out completely. So the net dipole moment is zero.
In SCl₂, sulfur has two lone pairs as well as two bonding pairs. Its electron geometry is tetrahedral, but its molecular shape is bent (V-shaped), not linear with two lone pairs on S
Because the molecule is bent, the two S–Cl bond dipoles do not cancel. So SCl₂ has a nonzero dipole moment.
ΔEN tells you whether an individual bond is polar, but it does not by itself tell you whether the whole molecule is polar.
To determine molecular polarity, you must draw the Lewis structure, find the shape of the molecule, and check whether the bond dipoles cancel due to symmetry.