Hydrogen bonding is a special type of dipole-dipole interaction that occurs when a hydrogen atom is covalently bonded to a highly electronegative, small atom (F, O, or N) and is attracted to a lone pair on another electronegative atom in a neighbouring molecule. In HF, the large electronegativity difference between H (2.1) and F (4.0) creates a highly polar H-F bond. The fluorine atom has three lone pairs and acts as a hydrogen bond acceptor, while the H atom acts as a hydrogen bond donor. This intermolecular hydrogen bonding creates an associated liquid state in HF, requiring extra energy to break these H-bonds upon boiling, hence the anomalously high boiling point (~19.5 degrees C) compared to HCl (-85 degrees C). London dispersion forces exist in all molecules but are generally weaker and increase with molecular size/surface area; they cannot explain why the lighter HF boils higher than heavier HCl. Dipole-dipole interactions are present in HCl as well, but HCl cannot form hydrogen bonds since Cl is not small and electronegative enough to support strong H-bonding. Ionic bonding is absent since HF is a covalent molecule. This is a direct application of NCERT Class 11 Chapter 4 (Chemical Bonding) on intermolecular forces and hydrogen bonding. Plausibility: the trend is reversed for HF compared to HCl, HBr, HI, which clearly signals a special interaction, confirmed as hydrogen bonding.