Why Aren't Chlorides Of The Noble Gases As Prevalent As Their Fluorides? I can't find the answer to this question on this SE website, and I apologize for the repetition if it has been answered before. 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?
-John
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).