Hydrogen bonding fundamentally shapes water's molecular structure through electrostatic attraction between the partially positive hydrogen atoms of one molecule and the partially negative oxygen of another [1][2]. This arises because oxygen's high electronegativity pulls shared electrons closer, creating the polarity necessary for these intermolecular bonds to form.

In liquid water, each molecule can form up to four hydrogen bonds—two as a donor through its hydrogens and two as an acceptor through its oxygen lone pairs [3][2]. However, these bonds are dynamic and weak (typically 10–40 kJ/mol), constantly breaking and reforming on femtosecond timescales [3][4]. This creates a fluctuating three-dimensional network rather than a fixed geometric arrangement, with no two molecules having identical hydrogen bond sets at any given moment [3].

The structural consequences are profound. Hydrogen bonding creates a tetrahedral local geometry that makes water less dense than it would be if packed like typical nonpolar liquids [4][5]. This open structure is responsible for water's anomalously low density in ice, causing it to float [5][6]. The bonds are strongest when nearly linear (O-H···O angle near 180°), and this directional preference influences crystal formation [4].

Hydrogen bonding also generates water's characteristic cohesion and high surface tension, as molecules at the surface are pulled inward by bonds from molecules below [2][7]. The network enhances molecular polarity by redistributing electron density, affecting water's chemical reactivity [5]. Additionally, quantum effects in hydrogen nuclei (particularly the light mass of protium) weaken bonds by approximately 10% compared to heavy water, demonstrating that even nuclear properties influence hydrogen bonding strength [3].

Without hydrogen bonding, water would be a gas at room temperature and lack the structural integrity that makes it Earth's universal solvent.