A hydrogen bond is a very strong type of dipole-dipole interaction. It occurs when a hydrogen atom is directly attached to nitrogen, oxygen, or fluorine.

Hydrogen bonds are stronger than typical dipole-dipole forces for two main reasons: the bonds are highly polarized due to a large electronegativity difference (e.g., in HF compared to CO), and hydrogen is a very small atom, leading to closer proximity and stronger electrostatic attraction.

Coulomb's Law dictates that electrostatic force is inversely proportional to the square of the distance between charges. Because hydrogen is small, the distance between nuclei is reduced, significantly increasing the attractive forces between molecules.

The bond within an HF molecule (covalent bond) is an intramolecular force. A hydrogen bond, however, exists between molecules (intermolecular force), such as the attraction between the fluorine of one HF molecule and the hydrogen of another.

In water, the partially negative oxygen atom of one molecule is attracted to the partially positive hydrogen atom of another, forming a hydrogen bond. This is an intermolecular force, distinct from the intramolecular covalent bonds within a single water molecule.

Vaporizing liquid water into steam requires breaking hydrogen bonds, an endothermic process (40.7 kJ/mol). In contrast, breaking an O-H covalent bond within a water molecule requires significantly more energy (467 kJ/mol). A covalent O-H bond is approximately 23 times stronger than the hydrogen bond between water molecules.