It is my understanding that compound formation has only been observed for the noble gases argon, xenon, and krypton, while no compounds have ever been observed for neon and helium, and that this might have something to do with both the decreased ionization potential of larger noble gas atoms and the high value of electronegativity of the halogen (usually fluorine) that bonds with the noble gas atoms to form compounds. So, my question is why aren't compounds formed from argon, xenon, and krypton bonding with chlorine, bromine, and iodine as prevalent as those formed from those noble gases and fluorine? Could the smaller size of the fluorine atom be a reason for its apparently higher reactivity?
You know that the bigger the noble gas atom, the easier it is to ionize it. By the same token it is also easier for a Xe atom than for a He atom to share some electron density in a covalent bond. That is, Xe is more reactive (a very little bit) than He (not at all).
For halogens, the tendency is in some sense the opposite: The smaller the halogen atom the more eager it is to gain electron density -- that is, the more electronegative it is. Fluorine is so electronegative that it can persuade Xe, Kr and even Ar (with difficulty) to give up some electron density to form a covalent bond. Cl is not quite so electronegative, and Br and I even less. Basically you need the least "noble" of the noble gases and the greediest halogen of all to even have a chance at a bond.
One can justify this handwaving with thermodynamic calculations and molecular orbital arguments (though the latter are often handwavy too).
Don't give up completely on noble gas chlorides. Xenon doesn't.
Xenon dichloride aside, it is true that noble gases are more apt to combine with fluorine. In that sense they are not so different from other nonmetals that form "expanded octet" compounds preferentially with fluorine.
Most readers are familiar with sulfur hexafluoride and its apparent twelve valence electrons in an "expanded octet" on the sulfur atom. In reality, of course, the four additional electrons do not actually overlap the sulfur atom; they are in ligand-based orbitals that are orthogonal to the sulfur valence orbitals. We call these orbitals "nonbonding" with their nonoverlapping the central sulfur atom, but they actually are antibonding between the ligands. Sulfur hexafluoride is stabilized by having relatively small ligands that hold their electrons tightly around a larger central atom; if we were to try to replace the fluorine atoms with larger and more polarizable chlorine atoms we'd aggravate the ligand-ligand antibonding as well as introduce unfavorable steric effects.
So it is with a noble gas as the proposed central atom, except that since the uncombined atom already has the closed-shell $ns^2np^6$ valence-electron configuration, just one pair of halogen atoms sets off the "expanded octet" condition with its potentially antibonding ligands. Xenon dichloride (which is linear) has a big enough central atom and the chlorine atoms far enough apart to get by (or, accordig to the Wikipedia article,it may be merely a van der Waals complex between a xenon atom and a dichlorine molecule), but putting those rather large chlorine atoms closer together around a noble-gas core is asking for trouble.