Hydrogen Bonding

Hydrogen bonding needs a hydrogen atom covalently bonded to nitrogen, oxygen or fluorine, and a lone pair on a nearby N, O or F. Those three elements are small and highly electronegative, so the hydrogen is left almost bare and can approach a lone pair very closely.

For the common molecules in this course, identify an N–H, O–H or F–H donor bond and an available lone pair on a suitable acceptor. The covalent O–H bond lies inside water; the hydrogen bond is the attraction from that H toward a neighboring acceptor atom, often drawn O–H···O. The H atom is not a free proton hopping between molecules in this picture. The N/O/F rule is an introductory rule for strong conventional hydrogen bonding, not a universal definition excluding every weaker interaction studied in advanced chemistry.

Having nitrogen or oxygen is not enough: its lone pair must be available. For example, an amide nitrogen's lone pair is delocalized into a carbonyl and is not treated as a usual hydrogen-bond acceptor. You need both a donor and an acceptor partner for the interaction. Acetone can accept from water even though pure acetone lacks the donors needed for conventional self hydrogen bonding.

Across group 16 the hydrides H₂S, H₂Se and H₂Te rise steadily with size, but H₂O sits far above the extrapolated line. NH₃ and HF show the same jump in their groups. Hydrogen bonding is the only difference.

At the same external pressure, a liquid boils when its vapor pressure reaches that pressure. The extensive hydrogen-bonding interactions in water keep its vapor pressure lower at a given temperature than a simple size-only comparison suggests. Water therefore must reach a higher temperature before boiling. The comparison changes intermolecular separation, not the energy required to break the covalent O–H bonds.

Each water molecule can form up to four hydrogen bonds. In ice those bonds lock molecules into an open hexagonal lattice that is less dense than liquid water, so ice floats. Hydrogen bonding also explains water high surface tension and its ability to dissolve alcohols and sugars.

Each water can donate through two H atoms and accept through its lone-pair directions, supporting a roughly tetrahedral network. In ice the ordered open arrangement leaves more space per molecule; the liquid retains hydrogen bonding but has a more disordered, partly collapsed network. Melting does not remove every hydrogen bond. Ice's lower density follows from this structure, rather than from water molecules individually expanding.

Short alcohols mix well with water because their O–H groups allow favorable water–solute hydrogen bonds. As the nonpolar carbon chain becomes longer, that region becomes a larger part of the molecule and aqueous solubility generally falls. A single O–H group is therefore not a guarantee of unlimited water solubility. Sugars have several hydroxyl groups that can interact with water.

For a property argument, compare hydrogen bonding in the pure liquids when discussing boiling, and compare solute–water interactions when discussing aqueous solubility. Those are different sets of partners. Boiling does not require eliminating every attractive interaction, and a numerical “strength of a hydrogen bond” is an approximate range rather than one constant for all donors and acceptors.