hybridization - Nitric acid vs. chloric acid strength

Here the central chlorine atom is best described as $\ce{sp^3}$ hybridized. This makes sense; the molecule is pyramidal. Also, due to the molecule's pyramidal geometry, this rules out the possibility of effective resonance due to less overlap between the $\ce{sp^3}$ lobes, which are angled, as opposed to the parallel pi lobes.

This explains quite well the difference in acidity between nitric acid and chloric acid; nitric acid is stronger than chloric acid. At first it seems difficult to differentiate between the acidities of both; the nitrate anion also has three peripheral oxygen atoms, which provide a similar (negative) inductive effect by delocalizing electron density.

With regard to electronegativity, the chlorate anion seems like it should be more slightly stable than the nitrate anion, implying that chloric acid is stronger than nitric acid; chlorine is more electronegative than nitrogen (if only slightly so - difference is 0.12 on the Pauling scale).

Oxidation state analysis suggests that the two should be similar in acidity; both the chlorine and the nitrogen have oxidation states of +5.

So it seems that the best rationalization of why nitric acid is stronger than chloric acid is the "better" (overlapping pi orbitals) resonance form of the nitrate anion.

With respect to the above rationalization,

1) How, exactly, do parallel pi orbitals enable "better" resonance? I get the principle of maximum overlap. If I had to guess, I would say that more overlap enables greater distribution of electrons among the relevant orbitals ... is there a more concise and scientific way of putting this? Is this even correct?

2) Are parallel pi orbitals necessary for resonance? I got the impression it is from organic class. Or are parallel pi orbitals only necessary for a special form of resonance - conjugation? Is the picture on the bottom referring to conjugation?

3) Are there any other ways to rationalize the difference in strength between these two acids?

Answer

The two ions you are comparing to each other are very different. In $\ce{ClO3-}$ there are two more electrons than in $\ce{NO3-}$. This is also the reason why the former pyramidalizes, while the latter does not.

In nitrate the negative charge can be delocalised over the whole molecule, since there is one molecular orbital with $\pi$ symmetry. One basic requirement for these orbitals is, that they are antisymmetric with respect to a mirror plane, that includes the bonding axis.

This MO scheme would also be thinkable for the chlorate ion, but here you have two more electrons, so you have to fill an antibonding orbital. The charge can no longer be delocalised as well in this molecule. The conclusion is, that the symmetry of the whole molecule is lowered and the $\pi$ orbitals vanish. A lone pair is always more stable in an orbital with high $\ce{s}$ contribution, which leaves only the $\ce{p}$ orbitals of chlorine for bonding with oxygen.

Resonance is a model which is necessary to describe bonding situations, that cannot be described by one Lewis structure alone. It is therefore more a rule of thumb than any evidence. $\pi$ bonds (or orbitals) are not necessary for valid resonance structures. They are however necessary for ([negative] hyper-)conjugation.