Strength of bases and acids Is it possible that strong acids like $\ce{H2SO4}$ behave as weaker ones like $\ce{CH3COOH}$ when the concentration is really high and vice versa for weak acid? And the same question for bases.
By this I am generally referring to pH values.

Why do concentrated acids react more than dilute ones? The degree of ionization of concentrated acids are lower than diluted acids, so they are supposed to react less. Where am I wrong?
In the case of weak acids dissociating like strong ones when they are highly dilute, there is a limit to how far this can go.  If you were to try to measure acetic acid dissociation at $10^{-n}$ molar for $n=4,5,6,7,8,...$ at ambient temperature you would at first rise nicely towards $100$% as you might expect, but eventually level off at only about $99.4$%.
What stops you is the autoionization of water.  As you approach infinite dilution the hydrogen ion concentration approaches not zero but the autoionized value of $10^{-7}$ molar.  So the ratio of dissociated acid to undissociated acid levels off at $K_a/10^{-7}$.  Accordingly the limiting degree of dissociation is
$1/[1+(10^{-7}/K_a)]=K_a/(K_a+10^{-7})$
For acetic acid with $K_a\approx1.8×10^{-5}$ this limit is about $99.4$%, but for a much weaker acid such as a typical phenol or hydro-cyanic acid the dissociation is much lower even at extremely low concentrations.  When $K_a&lt;10^{-7}$ the limiting dissociation is less than $50$%; you need a stronger base than water to coax dissociation as a major reaction in aqueous solution.