A hydrogen bond is an attractive force involving a hydrogen atom covalently bonded to a highly electronegative atom, such as nitrogen (N), oxygen (O), or fluorine (F). These electronegative atoms pull electrons toward themselves, causing the hydrogen atom to develop a large partial positive charge. This partially positive hydrogen is then attracted to a lone pair of electrons on a nearby electronegative atom. This attraction is not a true covalent bond, but a powerful, localized dipole-dipole interaction.
Defining Intermolecular and Intramolecular Forces
Intramolecular forces are the strong chemical bonds that exist within a single molecule, holding its atoms together. These include covalent bonds (where electrons are shared) and ionic bonds (the electrostatic attraction between oppositely charged ions). Breaking these forces requires significant energy and fundamentally changes the substance’s chemical identity.
Conversely, intermolecular forces are the weaker attractions that occur between separate molecules. These forces include London dispersion forces and dipole-dipole interactions, holding bulk matter together in its liquid or solid state. Overcoming these weaker forces causes a physical change, such as boiling or melting, but does not alter the molecular structure itself. Hydrogen bonds are generally classified as a strong type of intermolecular force.
The Role of Hydrogen Bonds Between Separate Molecules
Hydrogen bonds are most commonly recognized for their role as intermolecular forces, acting between separate molecules and significantly influencing physical properties. Water (H₂O) provides the most well-known example, where the oxygen atom of one molecule attracts the hydrogen atom of a neighboring molecule. Each water molecule can participate in up to four such bonds, creating a dense, interconnected network.
This extensive bonding explains the high boiling point of water compared to similar compounds like hydrogen sulfide (H₂S). A large amount of thermal energy is needed to break these attractions before the molecules can escape into the gas phase. Hydrogen bonds also contribute to water’s high surface tension and its density anomaly, where solid ice is less dense than liquid water because the bonds lock molecules into an open, crystalline lattice.
Other small molecules, such as ammonia (NH₃) and alcohols (containing an O-H group), also exhibit this bonding. Alcohols like ethanol have higher boiling points than organic molecules of similar size lacking the O-H group, due to these strong intermolecular interactions. The strength of a typical intermolecular hydrogen bond in a water-water interaction is approximately 21 kilojoules per mole.
Hydrogen Bonds Within a Single Molecule
Hydrogen bonds can also form within a single large molecule, known as intramolecular hydrogen bonding. This occurs when the hydrogen donor and acceptor atoms are part of the same molecule and the shape allows them to interact closely. This internal bonding is important in biology, as it determines the three-dimensional structures of macromolecules.
In proteins, intramolecular hydrogen bonds stabilize secondary structures, such as the alpha helix and the beta sheet. The alpha helix is held together by bonds forming between the hydrogen atom of an amide group and the oxygen atom of a carbonyl group four amino acids away along the same protein backbone. This regular pattern of internal attraction gives the helix its rigid, coiled structure.
In deoxyribonucleic acid (DNA), intramolecular hydrogen bonds hold the two complementary strands together in the double helix. These bonds form specifically between the nitrogenous bases: two bonds link Adenine (A) to Thymine (T), and three bonds link Guanine (G) to Cytosine (C). This internal attraction maintains the DNA’s functional shape and allows the strands to “unzip” during processes like replication and transcription.