Why is the solution of ammonium bifluoride more acidic than aqueous ammonium fluoride?
This question and the one referenced by Mithoron address NH$_4$HF$_2$ as an entity in itself which must be examined as a whole.
It is easier conceptually to rewrite the formula for the compound as NH$_4$F.HF. Consider adding NH$_4$F to water: you get a pH near 7. In a separate container, add HF to water (~0.1 M); you get a pH ~1 https://ehs.princeton.edu/book/export/html/197.
Combine the solutions and the pH will settle out somewhere between 1 and 7, because of the common ion effect: F$^-$ from the NH$_4$F inhibits the ionization of HF, but the solution is still acidic. The approximate pH can be calculated from the ionization constant of HF and the concentrations of NH$_4$F and HF.
The concept of the HF$_2^-$ moiety dominating the discussion for a solution is probably making the analysis murkier than it needs to be. It's just a hydrogen-bonded complex that happens to retain some semblance of existence in the solid. It is fairly stable: mp 126$^0$ C, bp 240$^0$ C (with decomposition). https://en.wikipedia.org/wiki/Ammonium_bifluoride