Homoassociation and Acid Strength
I think this Wikipedia article on hydrogen fluoride will answer your questions.  It discusses the acidity of $\ce{HF}$ from dilute to concentrated aqueous solution, and then pure $\ce{HF}$.  They present four equilibria to describe these situations

\begin{aligned}
\ce{H2O + HF &amp;~&lt;=&gt; [H3O^+ {.} F^{-}]&amp;&amp;(I)}\\
\ce{[H3O^+ {.} F^{-}] &amp;~&lt;=&gt; H3O^+ + F^{-}&amp;&amp;(II)}\\
\ce{[H3O^+ {.} F^{-}] + HF &amp;~&lt;=&gt; H3O^+ + HF2^{-}&amp;&amp;(III)}\\
\ce{3HF &amp;~&lt;=&gt; H2F^+ + HF2^{-}&amp;&amp;(IV)}\\
\end{aligned}

In dilute solution equilibrium (I) lies far to the right and the second equibrium (II) far to the left.  So even though the $\ce{HF}$ is extensively dissociated, the resultant hydrogen bonded ion pair is very stable and reluctant to undergo significant dissociation.  Hence the solution is weakly acidic.  As the concentration of $\ce{HF}$ is increased homodissociation comes into play [equilibrium (III)] producing the hydrogen difluoride ion as a discrete entity which is stabilized through intermolecular hydrogen bonding.  The increase in $\ce{H3O+}$ makes the resultant solution very acidic.  In pure $\ce{HF}$, autoionization occurs producing an exteremly acidic solution.  Page 4 of this link shows a graph of HF acidity from near zero to 50% concentration.