Again, I argue the problem arises by trying to explain why, when fluorine bonds are broken homolytically, they are much stronger than expected from the contributions of fluorine itself. It has been suggested by Pauling that bonds have covalent and ionic character and that it is fluorine's ionic character that increases its bond strength.
I do not agree with this. I do not believe this has been proven in any manner. If fluoride had high ionic content in a compound like HF, then its resistance to ionization virtually disproves itself. Pauling's calculations of electronegativity are not consistent with an additive principle of covalent and ionic content. He argues that opposite ionic content can only increase bond strength, yet it is well known that metal hydrides are weaker.
Except for the haloacetic acids, virtually all examples of electron withdrawing properties, acidity, reactivity, etc. point to the well known order I>Br>Cl>>F. Because of the Pauling electronegativity table, there are all sorts of rationalizations of why common sense fails to hold and how fluorine is not be a potent electron withdrawing group. To wit, the charge is more concentrated in fluorine, therefore it is a weaker acid. The bulk of atomic volume is the space occupied by the electrons. A low electron density (mostly empty) results in a low density atom. If the electrons are packed into a tighter space, that is more electrons per unit volume, it will have a higher electron density. If an electron's volume were constant, the density of atoms were vary with the neutron/proton ratio, or not very much. The reality is the nuclear charge of iodine is much greater than fluorine, its electrons are concentrated in a smaller volume per electron, it is a dense solid. Fluorine is a gas. If you took a measure, you would find more electrons in a smaller space around iodine. If electrons were negative, then the actual charge around iodine is larger. That isn't the entire story however. Iodine's nuclear proton field is larger and can repel a proton as well. The acidity is really a function of both fields, the local electron pair and the nuclear proton field. The closer iodine's electrons are to its nucleus, the closer and the more strongly will a proton feel the effects of the nuclear proton field as well.
I argue acidity is an inverse square result. Iodine, with more protons, pulls its electrons in more. Fluorine, with fewer protons does not. They reach out further. They react like a boxer, the one with the longer reach will hit you first. If you apply this simple principle, then HF, with its shorter bonds and electrons closer to its nucleus, is more acidic than H2O, and in turn more acidic than NH3 and more acidic than CH4. CH4 has the longest bonds. It has the fewest protons in its nucleus.
Acidity is a measure of heterolytic bond strength or about electron withdrawing properties. Electronegativity has been derived from bond strength data. It is actually a measure of homolytic bond strength. It is unfortunate that electronegativity arguments are as pervasive as they are. In my class, I simply divided electron withdrawing properties from electronegativity properties. You use electron withdrawing properties in the lab and electronegativity on the exam.
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Off topic
The people I would like to be critical of are the reviewers of Pauling's 1932 paper. We don't think about how something is thought to be true in science. Science is not a level playing field. Bigger reputations can get funding more easily than someone yet to prove themselves. A weak paper from a well regarded chemist can be accepted for publication more easily as well. However, to paraphrase myself, no matter how many times something is repeated, it will not be any more true than its original proof or proofs.
I think I have an explanation for the homolytic bond strength data for why fluorine bonds are stronger than expected. This is what Pauling was trying to achieve with electronegativity. However it becomes rather more complicated. It is in part a challenge to explain why the bonds of elemental fluorine as much weaker than expected. If you look at bond strengths, you will find the bonds are becoming stronger in the order I<Br<Cl and weaker in the order C>N>O. What is going to happen with fluorine and why? I just don't think it is because the electrons surrounding fluorine are more negative (ionic) than the electrons surrounding any other atom.